Solute and Solvent

In a binary solution, solvent is the component which is present in large quantity while the other component is known as solute.

Classification of Solutions

(A) Following types of solutions are seen on the basis of physical state of solute and solvent.

[if water is used as a solvent, the solution is called aqueous solution and if not, the solution is called non-aqueous solution.]

(B) Depending upon the amount of solute dissolved in a solvent we have the following types of solutions:

(i) Unsaturated solution A solution in which more solute can be dissolved without raising temperature is called an unsaturated solution.

(ii) Saturated solution A solution in which no solute can be dissolved further at a given temperature is called a saturated solution.

(iii) Supersaturated solution A solution which contains more solute than that would be necessary to saturate it at a given temperature is called a supersaturated solution.

Solubility

The maximum amount of a solute that can be dissolved in a given amount of solvent (generally 100 g) at a given temperature is termed as its solubility at that temperature.

The solubility of a solute in a liquid depends upon the following factors:

(i) Nature of the solute
(ii) Nature of the solvent
(iii) Temperature of the solution
(iv) Pressure (in case of gases)

Henry’s Law

The most commonly used form of Henry’s law states “the partial pressure (P) of the gas in vapour phase is proportional to the mole fraction (x) of the gas in the solution” and is expressed as

p = K_H . x

Greater the value of K_H, higher the solubility of the gas. The value of K_H decreases with increase in the temperature. Thus, aquatic species are more comfortable in cold water [more dissolved O₂] rather than Warm water.

Applications

1. In manufacture of soft drinks and soda water, CO₂ is passed at high pressure to increase its solubility.
2. To minimise the painful effects (bends) accompanying the decompression of deep sea divers. O₂ diluted with less soluble. He gas is used as breathing gas.
3. At high altitudes, the partial pressure of O₂ is less then that at the ground level. This leads to low concentrations of O₂ in the blood of climbers which causes ‘anoxia’.

Concentration of Solutions

The concentration of a solution is defined as the relative amount of solute present in a solution. On the basis of concentration of solution there are two types of solutions.

(i) Dilute solution
(ii) Concentrated solution

Methods of Expressing Concentration of Solutions

Various expression for the concentrations of solutions can be summarised as

(i) Percentage by weight (w / w %) It is defined as the amount of solute present in 100 g of solution.

w / w % = weight of solute / weight of solution * 100

(ii) Percentage by volume (w / V%) It is defined as the weight 01 solute present in 100 mL of solution.

w / V % = weight of solute / weight of solution * 100

or the volume of solute present in 100 mL of solution.

u / V % = volume of solute / volume of solution * 100

(iii) Mole fraction (x) It is defined as the ratio of the number of moles of a component to the total number of moles of all the components. For a binary solution, if the number of moles of A and B are n_A and n_B respectively, the mole fraction of A will be

(iv) Parts per million (ppm) It is defined as the parts of a component per million parts (10⁶) of the solution. It is widely used when a solute is present in trace quantities.

ppm = number of parts of the component / total number of parts of all the components * 10⁶

(v) Molarity (M) It is the number of moles of solute present in 1L(dm³) of the solution.

M = number of moles of solute / volume of solution (L)

M = mass of solute (in gram) * 1000 / mol. wt. of solute x volume of solution (in mL)

Molarity varies with temperature due to change in volume of solution.

[When molarity of a solution is 1 M, it is called a molar solution. 0.1 M solution is called a decimolar solution while 0.5 M solution is known as semi molar solution]

Molarity = Percent by mass * density * 10 / molecular weight

Dilution law, M₁ V₁ = M₂ V₂ (for dilution from volume V₁ to V₂)

For reaction between two reactants, M₁ V₁ / n₁ = M₂ V₂ / n₂

where, n₁ and n₂ arc stoichiometric coefficient in balanced equation.

(vi) Molality (m) It is the number of moles of solute per kilogram of the solvent.

Molality = mass of solute in gram * 1000 / mol. wt. of solute * mass of solvent (in g)

Molality is independent of temperature.

[Whcn solvent used is water, a molar (1 M) solution is more concentrated than a molal (1 M) solution.]

(vii) Normality (N) The number of gram equivalents of solute present in 1 L of solution.

Normality = number of grams – equivalent of solute / volume of solution in L

Number of gram-equivalents of solute = mass of solute in gram / equivalent weight

[Relationship between normality and molarity N x Eq. weight = M x mol. weight ]

If two solutions of the same solute having volumes and molarities V₁, M₁ and V₂, M₂ are mixed, the molarity of the resulting solution is

To dilute V₁ mL of a solution having molarity M₁ to molarity M₂ up to the final volume V₂ mL, the volume of water added is

(viii) Formality (F) It is the number of formula weights of solute present per litre of the solution.

Formality = moles of substance added to solution / volume of solution (in L))

(ix) Mass fraction Mass fraction of any component in the solution is the mass of that component divided by the total mass of the solution.

Molality, mole fraction and mass fraction are preferred over molarity, normality, etc., because former involve weights which do not change with temperature.

(x) Demal (D) It represents one mole of solute present in 1L of solution at O°C.

Raoult’s Law

The Raoult’s law states “For a solution of two volatile liquids, the vapour pressure of each liquid in the solution is less than the respective vapour pressure of the pure liquids and the equilibrium partial vapour pressure of the liquid is directly proportional to its mole fraction.

For a solution containing two liquids A and B, the partial vapour pressure of liquid A is

The proportionality constant is obtained by considering the pure liquid when χ_A = 1 then k = P°_A, the vapour pressure of pure liquid, hence

Konowaloff Rule

At any fixed temperature, the vapour phase is always richer in the more volatile component as compared to the solution phase. In other words, mole fraction of the more volatile component is always greater in the vapour phase than in the solution phase.

The composition of vapour phase in equilibrium with the solution is determined by the partial pressure of components. If Y₁ and Y₂ are the
component 1 and 2 respectively in the vapour phase then. using Dalton’s law of partial pressure,

p₁ = y₁ * P_total

p₂ = y₂ * P_total

Ideal Solutions

Those solutions in which solute-solute (B-B) and solvent-solvent (A-A) interactions are almost similar to solvent solute (A-B) interactions are called ideal solutions. These solutions satisfy the following conditions :

(i) Solution must obey Raoult’s law, i.e.,

(ii) ΔHmix = 0 (No energy evolved or absorbed)

(iii) ΔVmix = 0 (No expansion or contraction on mixing)

Some solutions behave like nearly ideal solutions, e.g., benzene + toluene. n-hexane + n-heptane, ethyl iodide + ethyl bromide, chlorobenzene + bromobenzene.

Non-ideal Solutions

Those solutions which shows deviation from Raoult’s law is called non-ideal solution.

For such solutions,

ΔH_mix ≠ 0

ΔV_mix ≠ 0

(a) Non-ideal solutions showing positive deviation In such a case, the A – B interactions are weaker than A – A or B – B interactions and the observed vapour pressure of each component and the total vapour pressure are greater than that predicted by Raoult’s law.

For such solutions

(b) Non-ideal solution showing negative deviation In such a case, the A – B interactions are stronger than A – A or B – B interactions and the observed vapour pressure of each component and the total vapour pressure are lesser than that predicted by Raoult’s law.

Azeotropic Mixture

A mixture of two liquids which boils at a particular temperature like a pure liquid and distils over in the same composition is known as constant boiling mixtures. These are formed by non-ideal solutions.

(i) Minimum boiling azeotropes are formed by those liquid pairs which show positive deviation from ideal behaviour. Such azeotropes have boiling points lower than either of the components, e.g., C₂H₅OH (95.57%) + H₂O (4.43%)(by mass).

(ii) Maximum boiling azeotropes are formed by those liquid pain; which show negative deviation from ideal behaviour. Such azeotropes have boiling points higher than either of the components. e.g., H₂O(20.22O%)+ HCl (79.78%] by mass.

Colligative Properties

[Colligatil1e : from Latin. = Co mean ‘together’; ligare means ‘to bind’.]

Colligative properties are those properties which depends only upon the number of solute particles in a solution irrespective of their nature.

Relative Lowering of Vapour Pressure

It is the ratio of lowering in vapour pressure to vapour pressure of pure solvent. The relative lowering in vapour pressure of solution containing a nonvolatile solute is equal to the mole fraction of solute in the solution.

Above expression is used to find the molecular weight of an unknown solute dissolved in a given solvent. Where, W_B and W_A = mass of Solute and solvent respectively. M_B and M_A = molecular weight of solute and solvent respectively.

Ostwald and Walker method is used to determine the relative lowering of vapour pressure.

Elevation in Boiling Point (ΔT_b)

Boiling point of a liquid is the temperature at which its vapour pressure becomes equal to the atmospheric pressure. As the vapour pressure of a solution containing a nonvolatile solute is lower than that of the pure solvent, it boiling point will be higher than that of the pure solvent as shown in figure. The increase in boiling point is known as elevation in boiling point, ΔT_b

ΔT_b = T_b – T°_b

ΔT_b = K_b m (where; m = molality)

K_b is molal elevation constant or ebullioscopic constant. Molecular mass of solute can be calculated as

where, W_B and W_A = mass of solute and solvent respectively.

K_b has units of K / m or K kg mol^-1, for water, K_b = 0.52 K kg mol^-1

The boiling point elevation of a solution is determined by

(i) Landsberger’s method

(ii) Cottrell’s method

Depression in Freezing Point (ΔT_f)

Freezing point of a liquid is the temperature at which vapour pressure of the solvent in its liquid and solid phase become equal. As we know that vapour pressure of solution containing non-volatile solute is lower than that of pure solvent, solid form gets separated out at a lower temperature as shown in the figure.

This decrease in freezing point of a liquid is known as depression in freezing point.

Depression in freezing point (ΔT_f) = T°_f – T_f

To find molecular mass of solute,

where, K_f is molal depression constant or cryoscopic constant.

K_f has units of K / m or K kg mol^-1.

Ethylene glycol is usually added to water in the radiator to lower its freezing point. It is called antifreeze solution.

[Common salt (NaCI) and anhydrous CaC12 are used to clear snow on the roads because they depress the freezing point of water. The freezing point depression is determined by Beckmann method or Rast method.]

Calculations of molal elevation constant (K_b) and molal depression constant (K_f)

Osmotic Pressure (π)

Osmosis is the phenomenon of spontaneous flow of the solvent molecules through a semipermeable membrane from pure solvent to solution or from a dilute solution to concentrated solution. It was first observed by Abbe Nollet.

Some natural semipermeable membranes are animal bladder, cell membrane etc.

CU₂[Fe(CN)₆]is an artificial semipermeable membrane which does not work in non-aqueous solutions as it dissolves in them.

Osmosis may be

(i) Exosmosis It is outward flow of water or solvent from a cell through semipermeable membrane.
(ii) Endosmosis It is inward flow of water or solvent from a cell through a semipermeable membrane.

The hydrostatic pressure developed on the solution which just prevents the osmosis of pure solvent into the solution through a semipermeable membrane is called osmotic pressure.

where, d = density, R = solution constant,

T = temperature, M_B = molar mass of solute

Osmotic pressure can be determined by anyone of the method listed below

(i) Pfeffer’s method
(ii) Berkeley and Hartley’s method (very good method)
(iii) Morse and Frazer’s method

On the basis of osmotic pressure, -the solution can be

(i) Hypertonic solution A solution is called hypertonic if its osmotic pressure is higher than that of the solution from which it is separated by a semipermeable membrane.

When a plant cell is placed in a hypertonic solution, the fluid from the plant cell comes out and cell shrinks, this phenomenon is called plasmolysis.

(ii) Hypotonic solution A solution is called hypotonic if its osmotic pressure is lower than that of the solution from which it is separated by a semipermeable membrane.

(iii) Isotonic solution Two solutions are called isotonic if they exert the same osmotic pressure. These solutions have same molar concentration. 0.91% solution of pure NaCl is isotonic with human RBC’s.

Two solutions are isotonic if they have the same molar concentration, e.g., if x % solution of X is isotonic with y % solution of Y, this means molar concentration of X = Molar concentration of Y

Osmotic pressure method is the best method for determining the molecular masses of polymers since observed value of any other colligative property is too small to be measured with reasonable accuracy.

Reverse osmosis When the external pressure applied on the solution is more than osmotic pressure, the solvent flows from the solution to the pure solvent, I which is called reverse osmosis. Desalination of sea water is done by reverse Osmosis.

Abnormal Molecular Masses

In some cases, observed colligative properties deviate from their normal calculated values due to association or dissociation of molecules. As we know,

Colligative property ∝ 1 / M_B

lienee, higher and lower values of molar mass is observed in case of association and dissociation respectively, e.g., in benzene, acetic acid gets associated, so, its observed molecular mass is 120. Similarly KCI undergoes dissociation in aqueous solution, so its observed molecular mass is 37.25.

These observed values are corrected by multiplying with van’t Hoff factor (i).

van’t Hoff Factor (i)

It is the ratio of observed value of colligative property to the calculated value of colligative property.

i = observed value of colligative property / calculated value of colligative property

or i = normal molecular mass / observed molecular mass

or i = number of particles after association or dissociation / number of particles initially

So to correct the observed value of molar mass, van’t Hoff factor (i) must be included in different expressions for colligative properties.

Degree of Dissociation (α) and van’t Hoff Factor (i)

(i) If one molecule of a substance gets dissociated into n particles or molecules and α is the degree of dissociation then

Degree of Association (α) and van’t Hoff Factor (i)

If n molecules of a substance A associate to form A_n and α is the degree of association then

van’t Hoff factor (i) > 1 for solutes undergoing dissociation and it is < 1 for solutes undergoing association.

Solids are the chemical substances which are characterised by define shape and volume, rigidity, high density, low compressibility. The constituent particles (atoms, molecules or ions) are closely packed and held together by strong interparticle forces

Types of Solids

The solids are of two types : Crystalline solids and amorphous solids.

Distinction Between Crystalline and Amorphous Solids

Structure Determination by X-ray Diffraction (Bragg’s Equation)

When a beam of X-rays falls on a crystal plane composed of regularly arranged atoms or ions, the X-rays are diffracted. If the waves are in phase after reflection, the difference in distance travelled by the two rays ti.e., path difference) must be equal to an integral number of Wavelength, nλ for constructive.

Thus, path difference = WY + YZ

= XY sin θ + xy sin θ

= 2 XY sin θ = 2d sin θ

∴ nλ = 2d sin θ

This equation is called Bragg’s equation.

Where, n = 1. 2, 3… (diffraction order)

λ = wavelength of X·rays incident on crystal

d = distance between atomic planes

θ = angle at which interference occurs.

Unit Cell

The smallest geometrical portion of the crystal lattice which can be used as repetitive unit to build up the whole crystal is called unit cell.

Types of Unit Cell

(i) Simple or primitive Unit cell In which the particles are present at the corners only.

(ii) Face centred unit cell In which the particles are present at the corners as well as at the centre of each of six faces

(iii) Body centred unit cell In which the particles are present at the corners as well as at the centre of the unit cell.

(iv) End centred unit cell In which the particles are present at the corners and at the centre of two opposite faces.

Number of Particles Per Unit Cell

Seven Crystal Systems

There are about 230 crystal forms, which have been grouped into 14 types of space lattices, called Bravais Lattices, on the basis of their symmetry and seven different crystal systems on the basis of interfacial angles and axes.

Seven Crystal Systems

Packing Fraction

It is defined as the ratio of the volume of the unit cell that is occupied by the spheres to the volume of the unit cell.

(i) Primitive cubic unit cell Atoms touch each other along edges.

Hence, d = a or r = a / 2

(r = radius of atom and a = edge length)

Therefore, PF = 4 / 3 πr³ / (2r)³ = 0.524 or 52.4%

(ii) Face centred cubic unit cell Atoms touch each other along the face diagonal.

Hence, d = a / √2

or r = √2a / 4

Therefore; PF = 4 * 4 / 3 πr³ / (4r / √2)r³ = 0.74 or 74%

(iii) Body centred cubic unit cell Atoms touch each other along the body diagonal.

Hence, √3a / 2

or r = √3a / 4

Therefore; PF = 2 * 4 / 3 πr³ / (4r / √3)r³ = 0.68 or 68%

Coordination Number

It is defined as the number of particles immediately adjacent to each particle in the crystal lattice.

[In simple cubic lattice, CN is 6, in body centred lattice, CN is 8 and in face centred cubic lattice, CN is 12].

High pressure increases CN and high temperature decreases the CN.

Close Packing in Crystals

Two Dimensional Packing of Constituent Particles

(i) Square close packing Space occupied by spheres is 52.4%.

(ii) Hexagonal close packing Space occupied by spheres is 60.4%.Hence. It is more efficient.

Three Dimensional Packing of Constituent Particles

(i) ABAB arrangement gives hexagonal close packing (hcp).

(ii) ABCABC arrangement gives cubic close packing or face centred CUbIC packing (ccp or fcc).

• In both these arrangements 740/0 space is occupied
• Coordination number in hop and ccp arrangement is 12 while in bcc arrangement, it is 8.
• Close packing of atoms in cubic structure = fcc > bcc > sc.
• All noble gases have ccp structure except He (hcp structure).

Void or Space or Holes

• Empty or vacant space present bet veen spheres of a unit cell, is called void or space or hole or interstitial void. When particles are closed packed resulting in either cpp or hcp structure, two types of voids are generated:
• Tetrahedral voids are holes or voids surrounded by four spheres Present at the corner of a tetrahedron. Coordination number of a tetrahedral void is 4.

• Octahedral voids are holes surrounded by six spheres located on a regular tetrahedron. Coordination number of octahedral void is 6.

[The number of octahedral voids present in a lattice is equal to the number of close packed particles. The number of tetrahedral voids present in a lattice is twice to the number of close packed particles.]

Density of Unit Cell (d)

Density of unit ce11 = mass of unit cell / volume of unit cell

d = Z * M / a³ = ZM / a³ * N_A

(The density of the unit cell is same as the density of the substance.)

where, d = density of unit cell

M = molecular weight

Z = no. of atoms per unit cell

N_A = Avogadro number

a = edge length of unit cell.

The Structure of Ionic Crystals

The ionic radius ratios of cation and anion, play a very important role in giving a clue to the nature of the crystal structure of ionic substances.

Radius Ratio and Crystal Structure

Ionic crystals may be of two types

(i)AB type and

(ii) A₂B or AB₂

Structure of Ionic Crystals

On applying pressure NaC} structure (6 : 6 coordination) changes into CsCI structure (8 : 8 coordination) and reverse of this occur at high temperature (760 K).

Imperfections in Solids

• In a crystalline solid, the atoms, ions and molecules are arranged in a Definite repeating pattern, but some defects may occur in the pattern. derivations from perfect arrangement may occur due to rapid cooling or presence of additional particles.
• The defects are of two types, namely point defects and line defects.

Point Defects

Point defects are the irregularities or deviations from ideal arrangement around a point or an atom in a crystalline substance Point defects can be classified into three types : (1) psychometric defects (2) impurity defects (3) non–stoichiometric defects

1. Stoichiometric Defect

These are point defects that do not disturb’ the -stoichiometric of the solid. They are also called intrinsic or thermodynamic defects. In ionic solids, basically these are of two types, Frankel defect and Schottky defect

AgBr has both Schottky and Frenkel defects. Frenkel defects are not found in pure alkali metal halides because cations are of large size.

2. Impurity Defect

• It arises when foreign atoms or ions aloe present in the lattice. In case of ionic compounds, the impurity 1S also ionic in nature. When the impurity has the same charge as the host ion. it just substitutes some of the host ions.
• Impurity defects can also be introduced by adding impurity ions having different charge than host ions. e.g. molten NaCl containing a little amount of SrCI₂ is crystallised. In such cases,
• Cationic vacancies produced = [number of cations of higher valence * Difference in valence of the host cation and cation of higher valence

3. Non-Stoichiometric Defect

Non-stoichiometric crystals are those which do not obey the law of constant proportions. The numbers of positive and negative ions present in such compounds are different from those expected from their ideal chemical formulae. However, the crystal as a whole in neutral.

Types of n-stoichiometric defects are as follows:

(i) Met excess defect Metal excess defect due to anionic vacancies: Alkyl halides like NaC1 and KCl show this type of defect. centres ale the sites from where anions are missing and the vacant sites are occupied by electrons. F-centres contribute colour and paramagnetic nature of the crystal [F stands for
German wo\d Farbe meaning colour).

Metal excess defect due to presence of extra cations at interstitial sites, e.g., zinc oxide is white in colour at room temperature. On beating, it loses oxygen and turns yellow.

(ii) Metal deficiency defect due to cation vacancy It is due to the absence of a metal ion from its lattice site and charge is balanced by ion having higher positive charge. Transition metals exhibit this defect, e.g., FeO, which is found in the composition range from Fe_0.93 O to Fe_0.96O.

In crystal of FeO, some Fe²⁺cations are missing and the loss of positive charge is made up by the presence of required number of Fe³⁺ ions.

Classification of Solids on the Basis of Electrical Conductivity

[The electricity produced on heating a polar crystal is called ‘pyroelectricity’.]

When mechanical stress is applied on polar crystals, electricity produced due to displacement of ions is called ‘piezoelectricity’

Semiconductors

Electronic conductors having electrical conductivity in the range of 10⁴ – 10⁷ Ω^-1 cm^-1 are known as semiconductors. Examples Si, Ge Sn (grey), Cu₂O, SiC and GaAs.

Intrinsic Semiconductors

Pure substances that are semiconductors are known as Intrinsic Semiconductors e.g., Si, Ge

Extrinsic Semiconductors

Their conductivity is due to the presence of impurities. They are formed by doping. It is defined as addition of impurities to a semiconductor to increase the conductivity. Doping of Si or Ge is carried out with P, As, Sb, B, Al or Ga.

(i) n·type semiconductors Silicon doped with 15 group elements like phosphorus is called n-type semiconductor. The conductivity is due to the presence of negative charge (electrons),

(ii) p·type semiconductors Silicon doped with 13 group element like gallium is called p-type semiconductor. The conductivity is due to the presence of positive holes.

• Some typical 13-15 compounds are InSb, AlP and GaAs and SOme typical 12-16 compounds are ZnS, CdS. CdSe and HgTe.
• These exhibit electrical and optical properties of great use in electronic industry.

Magnetic Properties of Solids

Solids can be divided into different classes depending on their response to magnetic field.

1. Diamagnetic Substances

These are weakly repelled by the magnetic field and do not have any unpaired electron, e.g., TiO₂, V₂O5, C₆H₆, NaCI, etc.

2. Paramagnetic Substances

These are attracted by the magnetic field and have unpaired electrons These lose magnetism in the absence of magnetic field, e.g., O₂, Cu²⁺, Fe³⁺, etc.

3. Ferromagnetic Substances

These are attracted by the magnetic field and show permanent magnetism even ill the absence of magnetic field e.g., Fe, Co and Ni.

4. Anti-ferromagnetic Substances

These substances have net magnetic moment zero due to compensatory alignment of magnetic . moments, e.g., MnO, MnO₂, FeO, etc.

5. Ferrimagnetic Substances

These substances have a net dipole moment due to unequal parallel and anti-parallel alignment of magnetic moments, e.g., Fe₃O₄, ferrites, etc.

The branch of chemistry which deals with the identification of constituents of a substance and calculation of their amounts is called analytical chemistry.

Branches of Analytical Chemistry

The branches of analytical chemistry are

(a) Qualitative analysis It deals with the identification of various constituents present in a given material. It further be classified as

1. Qualitative analysis of inorganic compounds
2. Qualitative analysis of organic compounds

(b) Quantitative analysis It deals with the measurement of amounts/volume/strength of a substance/solution.

Qualitative Analysis of Inorganic Compounds

It involves the following steps

1. Preliminary tests
2. Dry tests
3. Wet tests for anions
4. Wet tests for cations

Preliminary Tests

(a) Solubility Compounds that dissolve in water to the extent of approximately 0.02 mol per litre (0.02M) are usually classified as “soluble” compounds, while those that are less soluble are classified as “insoluble” compounds. No gaseous or solid substances are infinitely soluble in water.

1. All common inorganic acids are soluble in water. Low molecular weight organic acids are soluble.
2. All common compounds of the group IA metal (Na, K etc) and the ammonium ion, NH₄⁺, are soluble in water.

3. All common nitrates NO₃^– acetates, CH₃COO^– and perchlorates, CIO₄COO^–, are soluble in water.

4. (a) All common chlorides, CI^– , are soluble in water except AgCl, Hg₂Cl₂ and PbCl₂.

(b) The common bromides Br^– , and iodides I^–, show approximately the same solubility behaviour as chlorides. but there are some exceptions. As the halide ions (Cl^– , Br^–, I^–) increase in size, the solubilities of their slightly soluble
compounds decrease. Although HgCl₂ is readily soluble in water, HgBr₂ is only slightly soluble and HgBr₂ is even less soluble.

(c) The solubilities of the pseudo-halide ions, CN^– (cyanide) and SCN^– (thiocyanate), are quite similar to those of the
corresponding iodides. Additionally, both CN^– and SCN^– show strong tendencies to form soluble complex compounds.

5. All common sulphates, SO₄^2-, are soluble in water except those PBSO₄, and HgSO₄.CaSO₄and Ag₂SO₄ are sparingly soluble.

6. All common metal hydroxides are insoluble in water except those of the group IA metals and the lower members of the group IIA metals, beginning with Ca(OH)₂.

7. All common carbonates CO₃^2- phosphates, PO₄^3-, and arsenates AsO₄^3-, are insoluble in water except those of the group IA metals and NH₄⁺.MgCO₃ is fairly soluble.

8. All common sulphides, S^2- are insoluble in water except those of the group IA and group IIA metals and the ammonium ion.

(b) Colour change on beating Certain oxides change colour on heating and this fact can be used to identify salt.

Dry Tests

(a) Borax bead test If borax, Na₂B₄O₇·10H₂O is heated on the platinum loop, a transparent colourless glass like bead of sodium metaborate (NaBO₂) and boric anhydride B₂O₃) is formed.

Colour In Oxidising and Reducing Flames In Borax-Bead Test

(b) Microcosmic salt bead test A test similar to borax bead test is used for identification of coloured cations if microcosmic salt, Na(NH₄)HPO₄·4H₂O, is used.

Results have been summarisod in the following table.

Microcosmic Salt Bead Test

(c) Sodium carbonate bead test The sodium carbonate bead is prepared by fusing a small quantity of sodium carbonate on a platinum wire loop in the Bunsen flame; a white, opaque bead is produced. If this is moistened, dipped into a littie KNO₃ and then into a small quantity of a manganese salt (for example) and the whole heated in the oxidising flame, a green bead of sodium manganate (Na₂MnO₄) is formed.

(d) Flame test Paste of the salt and cone. HCl is taken into the lower oxidising zone and colour imparted to the flame by salts is observed. Salts, particularly of group V (Ba²⁺, Ca²⁺, Sr²⁺) are identified by colours of the flame and summarised in Table

Flame Tests

The yellow colouration due to sodium masks that of potassium. In such cases view the flame through cobalt glass, the yellow sodium colour is absorbed and the potassium flame appears crimson.

Wet Tests for Anions

Salt or mixture is treated with dil. H₂SO4 and also with cone, H₂SO4 separately and by observing the types of gases evolved, confirmatory tests of anions are performed.

Observation with Dilute H₂SO4

Observation with Concentrated H₂SO4

(i) Chloride

Chromyl-chloride test

(ii) Bromide

(iii) Iodide

Wet Tests for Cations

Different basic radicals have different solubility products and precipitated by different reagents. Thus, on the basis of solubility products and reagent used, basic radicals are classified into following groups.

Classification of Basic Radicals Into Groups Based on k values

(c) Group III

Boil off H₂S gas from the filtrate of group II. Add NH₄Cl and one drop of dil. HNO₃; heat, cool and add NH₄OH → ppt

• Reddish brown ppt of Fe(OH)₃ if Cr³⁺ is present (salt is brown).
• Dirty green ppt of Cr(OH)₃ if Cr³⁺ is present (salt is green).
• White ppt of Al(OH)³⁺ if Al³⁺ is present (salt is colourless).

(d) Group IV

White ppt of Zns reappears on passing H₂S into soluble Na₂ZnO₂ solution.

(ii) Mn²⁺ MnS (buff coloured ppt) are soluble in dil. HCl but addition of excess ofNaOH gives Mn(OH)2 ppt which changes to brown/black by atmospheric O₂.

Dissolve MnO₂ or Mn(OH)₂ ppt into HNO₃ and then add oxidising agent KClO₃ or PbO₂; heat to boil-purple solution (B).

Some Other Important Tests

(a) Potassium Iodide (KI)

(c) K₄ [Fe(CN)₆] Potassium hexacyanoferrate (II) (Potassium ferrocyanide)

Qualitative Analysis of Organic Compounds

It also involves following steps:

1. Preliminary test
2. Element detection
3. Test of functional group.

1. Preliminary Tests

(a) Heat test The compound is heated in flame. If it burns with smoky flame, it is an aromatic compound and if it burns with
non-smoky flame, it is an aliphatic compound.

(b) Colour

(c) Odour

Element Detection

See chapter purification and characterisation of organic compounds.

Tests of Functional Groups

(a) Tests for Carboxylic (-COOH) Group

(i) Litmus paper test Dip blue litmus paper in the aqueous solution or suspension of the compound. It turns red.

(ii) Sodium bicarbonate test In a test tube take a little quantity of the compound and then, add a saturated solution of sodium bicarbonate. Formation of brisk effervescence shows the presence of -COOH group.

(iii) Ester formation Heat a small quantity of organic compound with ethyl alcohol and a little conc H₂SO₄. Cool the solution and pour in a tube containing water. A fruity smell, due to formation of an ester, indicates the presence of carboxylic group.

(b) Tests for Alcoholic (-OH) Group

(i) Certc ammonium nitrate test To small amount of organic compound or its aqueous solution, add a few drops of ceric ammonium nitrate. A red colour indicates the presence of alcoholic hydroxy group.

(ii) Lucas test This test is used to di stinguish between primary, secondary and tertiary alcohols. In this test, treat 2 mL of organic compound with about 8 mL of Lucas reagent (for preparing Lucas reagent dissolve 32 g of anhy ZnCl₂ in 20 mL of cone HCl) and shake.

(a) Immediate formation of turbidity indicates the presence of tertiary alcohol.

(b) Formation of turbidity after 4-5 min shows the presence of secondary alcohols.

(c) If solution remains clear, then primary alcohol is present.

(c) Tests for Phenolic (Ph-OH) Group

(i) Ferric chloride test To aqueous or alcoholic solution of compound, add few drops of ferric chloride (FeCl a). Formation of green, blue or violet colour shows the presence of phenol.

(ii) Liebermann’s nitroso reaction Fuse a little amount of compound with a crystal of NaNO₂ in a test tube. Cool the mixture and add 1 mL conc H₂SO₄.A deep green to blue solution is formed which turns red when poured in a large excess of water. The red aqueous solution becomes again deep green or blue if made alkaline with NaOH. It shows the presence of phenol.

Note Nitrophenol does not respond to FeCl₃ test as well as Liebermann’s nitroso reaction.

(d) Tests for Aldehyde (-CHO) Group

(i) Tollen’s reagent test Take a little quantity of the compound in a test tube and add 2 mL of freshly prepared reagent. Shake, warm and allow the contents to stand for 2-3 min. Formation of silver mirror or a grey ppt indicates the presence of an aldehydic group.

(ii) Fehling’s solution test Take a mixture of equal amounts of Fehling’s solution A and B, and a few drops of organic compound and boil the contents. Formation of a red ppt shows the presence of an aldehyde.

Note : Both the above test» are also given by reducing sugars.

(iii) Schiff’s reagent test Add 5-6 drops of organic compound to 2 mL of the reagent. Shake vigorously. Mer some time formation of a deep red or violet colour indicates the presence of an aldehydic group.

(iv) Benedict’s solution test Boil the compound with 2-3 mL of Benedict’s solution for few minutes. Appearance of a red-yellow ppt confirms the presence of aliphatic aldehydes.

Note : This test is usually given by only aliphatic aldehydes, thus used to differentiate between aliphatic and aromatic aldehydes.

(i) 2,4-dinitro phenyl hydrazine test In a dry test tube add few drops of the organic compound (if liquid) or its alcoholic solution (if solid) to about 2 mL of the reagent and one drop of cone H₂SO₄. Shake vigorously, heat (if necessary) and allow to stand for about 5 min. A yellow or orange ppt separates out in case of a compound containing carbonyl group due to formation of respective hydra zones.

(ii) Sodium nitroprusside test Add 0.1 g of solid or 0.2 cc of liquid compound to 2 cc of sodium nitroprusside solution and then, make it alkaline with 2-3 drops of sodium hydroxide. A red or purple colour indicates the presence of ketone (benzophenone does not give this test).

(f) Tests for PrImary Amine (-NH₂) Group

(i) Carbylamine test Boil a little quantity of the compound with 2 drops of chloroform and 2 mL of alcoholic caustic potash.

An intolerable offensive odour of carbylamine indicates the presence of primary amine.

(ii) Dye test Dissolve about 0.2 g of the compound in dil HCl and cooL Now, add 10% aq NaNO₂solution. Pour all this content into a beaker containing alkaline β-naphthol solution. Formation of a red or orange dye indicates the presence of aromatic primary ammo group.

(g) Tests to Distinguish between Primary, Secondary and Tertiary Amlnes

(i) Nitrous acid test Prepare a solution of nitrous acid by adding ice cold dil HCl to a solution of 1% aq NaNO₂. Add gradually this solution to 0.2 g of the organic compound in 10 mL dil HCl.

(a) Formation of brisk effervescence shows the presence of aliphatic primary amine.

c) No reaction indicates the presence of aliphatic tertiary amine while production of green or brown colour indicates the presence of aromatic tertiary amines.

(ii) Hinsberg’s test To about 0.2 g of the compound) add 1 mL of 5% NaOH and 3 mL pyridine. Shake well and add few drops of benzene sulphonyl chloride with continuous shaking.

(a) Formation of yellow colour indicates the presence of primary amine.

(b) Formation of orange colour shows the presence of secondary amine.

(c) Formation of a red or purple colour shows the presence of tertiary amines.

Preparation of OrganiC Compounds

1. Acetanilide

Amines containing -NH₂ and > NH groups respectively can be directly acetylated. Their reactive hydrogen atoms get replaced by the acetyl group (-COCH₃) to give acetyl derivatives of the type RNH·COCH₃ and R2N . COCHa respectively which may be regarded as mono and di-alkyl substituted acetamide.

2. p-nitro Acetanilide

When acetanilide is treated with a mixture of cone HNO₃ and conC H₂SO₄, it gives p-nitro acetanilide alongwith a little amount of o-isomer. In this process fuming HN0 3 in the presence of conc H₂SO₄ gives nitronium ion (NO₄⁺) which attack on acetanilide to form p-nitro acetanilide through the cyclopentadienyl cation (intermediate) formation.

3. Aniline Yellow

When diazo amino benzene is heated with aniline and a little amount of aniline hydrochloride at about 40°C, for a short time, it gives p-amino azo benzene or aniline yellow in good yield.

The mechanism of the reaction is based on the equilibrium involving the diazo amino compound, phenyl diazonium chloride and aniline.

The reaction takes place between the two latter compounds under weakly acidic conditions.

4. Iodoform

Acetone when treated with potassium iodide and sodium hypochlorite (NaOCl), gives iodoform.

Titrimetric Exercises

Some Important terms

(a) Titration The process by which the concentration or strength of a chemical substance is measured by using the solution of known strength is called titration.

(b) Analyte and titrant The substance being analysed is called analyte and that which is added to analyte in a titration is called titrant.

(c) Equivalence point or end point It is point at which the reaction between two solutions is just complete. It is generally represented by change in colour, pH, conductivity etc.

(d) Standard solution Solution of known concentration is called standard solution.

(e) Primary standard substance The substance standard solution of which can be prepared directly by dissolving its definite weight in definite volume of solvent is called primary standard substance, e.g., crystallie oxalic acid, anhydrous Na₂CO3, Mohr’s salt etc.

The substance, which occur in pure state, are non-hydroscopic, non-deliquescent, generally behave as primary standard substance.

(f) Secondary standard substance Their standard solution cannot be prepared directly. e.g., KMnO₄, NaOH, KOH etc.

(g) Indicator It shows the end point of a titration.

Na₂CO₃ vs HCl Titration

The titration of Na₂CO₃ us HCl is a neutralisation titration (acidimetry and alkalimetry) which involve the neutralisation of an acid with a base. e.g., sodium carbonate is attacked by dil HCl in the following way

and then titrated its 20 mL with HCI solution by adding a few drops of methyl orange indicator. Change in colour shows the end point. Note the reading and repeat the procedure to obtain concurrent readings.

Volume of HCI Used with 20 mI, of Known (prepared) Na₂CO₃ Solution

Calculations

(i) Weight of Na₂CO₃ dissolved in 250 mL measuring flask = z

= ….. g

Weight of Na₂CO₃ in 100 mL = ….. x 4

= ……. g/L

Normality of Na₂CO₃ (prepared) = [strength(g/L)/Eq. wt. of oxalic acid)

= (…… N/53)

• In acidimetry and alkalimetry, the choice of indicators mainly depends upon the nature of the acids and alkalies used. Methyl orange, phenolphthalein are some of the important indicators used in these titrations.
• As no indicator gives correct results in the titration of weak acids against weak bases, such titrations are to be avoided.

Oxalic Acid vs KMnO₄ Titration

This is an example of redox titrations, in which a reducing agent (as oxalic acid) is estimated by titrating it with a standard solution of oxidising agent (as KMnO₄). Such reactions are accompanied by the change in valency of ions. In these titrations oxidation and reduction takes place simultaneously i.e., while one substance is being oxidised, the other one is being reduced.

The last drop of KMnO₄ itself acts as an indicator (self indicator).

In this titration, first the standard solution of oxalic acid is prepared which is than titrated KMnO₄solution in the presence of dil. H₂SO₄. The procedure is repeated to obtain a set of concurrent readings.

Calculations

(ii) For the titrations using standard oxalic acid solution

(iii) For the titration using supplied oxalic acid solution

Sulphuric acid should be in excess otherwise a brown ppt due to formation of MnO₂ will be formed.

This titration cannot be carried out in presence of acid like HNOs and HCI, because HNO₃ itself is an oxidising agent, so it will interfere with the oxidising action of KMnO₄ and HCl reacts chemically with KMnO₄ solution.

Mohr’s Salt n KMnO₄ Titration

This titration is also an example of redox titrations and work on the same principle as oxalic acid us KMnO₄ titration. In this titration the active constituent of ferrous ammonium SUlphate (Mohr’s salt) is

ferrous sulphate, which is oxidised to ferric sulphate by acidified potassium permanganate as follows.

Calculations

(ii) For the titrations using standard ferrous ammonium sulphate solution

(iii) For the titration using supplied ferrous ammonium sulphate solution

N₃V₃ = N₄V₄

ferrous ammonium sulphate KMnO₄ [because N₄ = N₂](unknown)

N₃ = …N

Strength of ferrous ammonium sulphate in gIL = N₃ x Eq. wt. of ferrous ammonium sulphate

= … g/L

Nucleons and Nuclear Forces

Protons and neutrons which reside in the nucleus, are called nucleons and forces binding them in the nucleus, are called nuclear forces.

These are short range forces operating over very small distances 1 fermi = 10^{– 15}cm).

These forces are 10²¹ times stronger than the electrostatic forces.

n^o and p⁺ are held together by very rapid exchange of nuclear particles, called mesons (discovered by Yukawa). Mesons may be positively charged (π⁺), negatively charged (π^–) or neutral (π^o).

Parameter of Nucleus

(i) Radius of nucleus It is proportional to the cube root of the mass number of element.

R = R_o * A^{1 / 3}

Where, A = mass number, R = radius of nucleus

R_o = proportionality constant = 1.4 x l0^{– 13} rn= 1.4 x 10^{– 15}cm

Radius of the nucleus = 10^{– 13} – 10^{– 12}cm

Factors Affecting Stability of Nucleus

Stability of nucleus is affected by following factors

1.Neutron-proton Ratio or n / p Ratio

It is the main factor for determining the stability of nucleus.

The nuclei located in the zone of stability or belt of stability are stable.

Nuclei lying above this zone, have higher n / p ratio and undergo β – emission to get the zone of stability.

Here, a n^o is transformed into p⁺ and give β and antineutrino. Thus, emission of β particles increases the number of proton and decreases the number of neutrons.

2. Binding energy

It is the energy released when a nucleus is formed from its constituent nucleons.

Higher the BE per nucleon, greater is the stability of the nucleus. Fe has maximum BE of 8.8 MeV per nucleon. Hence. it is most stable.

3. Packing fraction

It is the measure of relative mass defect.

Packing fraction = isotopic mass – mass no / mass no. * 10⁴

Its value may be positive, negative or zero. Nuclei with positive packing fraction are highly unstable. Packing fraction is least for Fe and highest for H₂.

4. Magic numbers

Nuclei having 2, 8, 20, 28, 50, 82 and 126 protons or neutrons are stable, hence these numbers are called magic numbers.

Nuclei with even numbers of both p⁺ and n^o are generally more stable than those with their odd numbers.

Radioactivity

Certain nuclei spontaneously emit radiations. Such nuclei are called radioactive and this phenomenon of disintegration of nuclei spontaneously is called radioactivity.

It was discovered by Henri Becquerel in 1895. However, term radioactivity was given by Madam Curie.

Marie Curie and her husband Piere Curie isolated polonium and radium from pitchblende (UaOs)’

Radioactivity being a natural nuclear phenomenon, remains unaffected by external factors like temperature, pressure etc. Potassium uranyl sulphate [K (UO₂) (SO₄)₂] was the first compound found to be radioactive. Tritium (₁H ) is the lightest radioactive element.

Radioactive Radiations

Radioactive radiations are of three types : Positively charged α-rays, negatively charged β-rays and uncharged γ-rays.

Properties of α, β and γ-rays

Radioactive Disintegration

The phenomenon of conversion of one radioactive nuclei into another with the emission of α, β and γ-rays is called radioactive disintegration.

It was proposed by Rutherford and Soddy in 1903.

γ-rays are secondary effects of radioactive disintegration i.e, emitted only after emission of α and β-rays.

Methods of Radioactive Disintegration

No. of β -particles emitted = 2 x number of α -particles – (change in atomic number)

Emission of one α and two β-particles gives isotope.

(iii) γ-decay Emission of γ-rays does not affect the atomic number as well as mass number.

Group Displacement Law

This law was given by Soddy, Russel and Fajans in 1913 to decide the position of element, obtained after radioactive disintegration, in the Periodic Table.

According to this law, ‘In an a-emission, the daughter nuclei will occupy a position which is two places left to the position of parent nuclei and in a p-emission, the daughter nuclei will occupy a position one place right to the parent nuclei”.

Rate of DiSintegration

It depends only upon the number of atoms present and is defined as the number of atoms of radioactive element that disintegrate in a unit time.

Half-life Period, t_1/2

It is defined as the time required for one-half of the isotope to decay. It is calculated as

t_1/2 = 0.693 / k

Its unit is t^-1.

Half-life is related to total time as, T = n x t_1/2

Average Life (λ)

The statistical average of the lives of all atoms present at any time is called the ‘average life’. It has been shown to be reciprocal of decay constant, k.

Activity of Radioactive Substance

Activity is defined as the number of disintegrations occurring in a radioactive substance per second.

Higher is the activity of a substance, faster will be its disintegration.

Geiger Muller counter is used to determine the activity of a radioactive sample.

Units of Radioactivity

(i) Curie The activity of a substance, disintegrating with a rate of 3.7 x 10¹⁰ disintegration per second (dps), is said to be 1 curie.

1 Curie = 3.7 * 10¹⁰ dps

1 millicurie = 3.7 * 10⁷ dps.

(ii) Rutherford It is the activity of a radioactive substance which disintegrates with a rate of 10⁶dps.

1 Rutherford (rd) = 10⁶ dps.

(iii) Becquerel It is the activity of the substance, which disintegrates with a rate of 1dps. It is the 81unit of radioactivity.

1Bq (Becquerel) = 1 dps

Disintegration Series

These are the decay series in which heavy nuclei decay by a series of α and β emissions finally resulting in the formation of stable isotopes of
lead.

There are three natural disintegration series while the fourth series, called the neptunium series is artificial

Radioactive Disintegration Series

To find a series, mass number is divided by 4.

Artificial Radioactivity

The phenomenon of conversion of a stable nuclei into a radioactive nuclei by bombarding it with a high speed projectile is called artificial radioactivity or induced radioactivity (given by Curie and Joliot).

The element with atomic number greater than 92 (after U) are obtained by this process. These are called transuranium elements.

These all are synthetic and radioactive

Artificial Transmutation

The conversion of a non-radioactive element into another non-radioactive elements by artificial means, i.e, by bombarding with some fundamental particles, is called artificial transmutation. e.g,

Nuclear Reactions

In these reactions, the nuclei of atoms interact with other nuclei or elementary particles such as α, β, p, d, n etc., and results in the formation of a new nucleus and one or more elementary particles.

A nuclear reaction can be represented as

parent nuclei (projectile, obtained particle) daughter nuclei

e.g, the reaction,

Energy of a nuclear reaction, Q = (total mass of products – total mass of reactants) x 931.5 MeV.

For exoergic reactions, Q is negative and for endoergic reactions, Q is positive.

Types of Nuclear Reactions

Accelerators

These are used to give projectiles. Commonly used accelerators are linear accelerator, cyclotron and synchrotron.

Synchrotron is used as proton accelerator. Neutron being neutral are best projectile, protons and deuterons being monopositive are better projectile than u-particles (dipositive).

Nuclear Fission (0 Hahn and Fritz)

It is the process of splitting of a heavy nucleus into fragments with the emission of a large amount of energy. e.g,

Nuclear fission is a chain reaction, e.g., it involves a series of reactions in which the product of the reaction initiate other similar reactions. e.g,

During fission, there is always lose of mass which is converted to energy according to Einstein equation.

E = Δmc²,(where, Δm = mass defect)

The minimum mass of fissionable material required to sustain a chain reaction is called critical mass. If the mass of the fissionable material is more than the critical mass, it is called supercritical mass and if the mass of the fissionable material is smaller than the critical mass, it is called subcritical mass.

For U-235, the critical mass lies between 1 kg to 100 kg.

A nuclear fission process may be controlled or uncontrolled. Reactors are the examples of controlled chain reaction while atom bomb is an example of uncontrolled chain reaction.

Nuclear Reactor

In it, the nuclear fission process is carried out in a controlled manner. control rods made up of B or Cd absorb additional neutrons and thus, slow down the reaction. A nuclear reactor consists of :

(a) Fissile material U enriched in ₉₂U²³⁵ is used as fissile material.
(b) Moderator These are used to slow down the speed of neutrons. e.g, graphite, heavy water, (D₂O).
(c) Control rods Rods of B or Cd.
(d) Coolant Liquid alloy of Na and K is used as coolant. It takes away the heat of the exchanger. Heavy water, polyphenols and CO₂ have also been used as coolants.

Breeder Reactor

This reactor also uses uranium as fuel but unlike a conventional reactor, it produces more fissionable materials than it uses.

U-235 or Pu-239 mixed with U-238 is used as a fuel in a typical Breeder reactor, so that breeding takes place within the core. Here, U-235 (or Pu-239) undergoes fission and U-238 capture more than one neutron (which are obtained during fission) to generate Pu-239. Thus, the amount of fissionable material can be steadily increased as the
starting nucleus fuels are consumed.

Fertile and fissile materials A fertile material is non-fissile in nature but can be made possible by reactions with neutrons. e.g, U²³⁸,Th²³².

Fissile materials readily undergo fission and produce a large amount of energy. e.g,U²³⁵, Pu²³⁹ , U²³³ etc.

Atom Bomb

It contains two pieces of a fissile material having subcritical mass and an explosive. At the time of explosion, the explosive explodes and brings the subcritical masses together. Thus, a single substance having supercritical mass is obtained, which starts chain reaction and a huge amount of energy is released leading to explosion.

Nuclear Fusion

It is the process of conversion of two or more lighter nuclei into a heavier nuclei with the emission of a large amount of energy.

For the initiation of these reactions, a very high temperature (> 10⁶K) is required so, these are also called thermonuclear reactions.

Energy of Sun and hydrogen bomb are based on the principle of nuclear fusion.

Following reactions are assumed as the source of energy of the Sun.

Every second, the Sun loses 4.3 X 10⁹ kg (430000 tons) of mass by the fusion process. But the total mass of the Sun is so great that this loss of mass is imperceptible.

Hydrogen Bomb

It contains a mixture of D₂O and T₂O (tritium oxide) surrounding an ordinary atomic bomb. The temperature produced by the explosion of the atomic bomb initiates the fusion reaction between ₁H³ and ₁H²

Applications of Radioactivity

1. Estimation of Age (Dating Technique)

(i) Carbon dating technique It is used to determine the age of wood, animal fossils etc. It is based upon the ratio of C¹⁴ to C¹² which remains constant in living organisms but decreases in dead sample. By comparing these two, the age is determined

2. Medicinal Use

Radioisotopes are used to diagnose and cure many diseases. These can be used in three ways.

(i) In Vivo studies ⁵¹Cr is used for such technique.

(ii) In therapeutic procedure (to cure diseases) Co- 60 is used for the treatment of cancer, Na-24 is injected to trace the Dow of blood, 1-131 is used for the treatment of thyroid and P-32 is used for leukemia.

(iii) Imaging procedure 1-131 is used to study the structure and activity of thyroid gland. 1-123 is used in brain imaging and Tc-99 M is used in bone scans.

Radioisotopes are also widely used to find the reaction mechanism, in industry and in agriculture.

Classification of Hydrocarbons

Alkanes

Alkanes are saturated, open chain hydrocarbons containing carbon-carbon single bonds. e.g., methane (CH₄), ethane (C₄H₆) propane (C₃H₈), etc.

These hydrocarbons are inert under normal conditions [i.e.,do not react with acids. bases and other reagents). Hence, they were earlier known 88 paraffins (Latin : parum-little; affins-affinity)

Alkanes exhibit chain isomensm, position isomerism and conformational isomerism.

Methods of Preparation of Alkanes

i) From hydrogenation of alkenes and alkynes

Ease of hydrogenation depends on the steric crowding across the multiple bond. More is the steric crowding, the less is the
reactivity towards hydrogenation.

(ii) By sodalime Decarboxylation of sodium or potassium salts of fatty acids [decarboxylation reaction]

This reaction is used for descending of series as the alkane obtained has one carbon less than the parent compound. CaO is
more hygroscopic than NaOH and it keeps NaOH in dry state.

(iii) By Wurtz reaction

This reaction is used to increase the length of the carbon chain.

(iv) By reduction of alkyl halides

Reducing agents like Zn/HCl, HI/Red P, H₂/Pd can also be used.

(v) By Kolbe’s electrolysis

Only alkanes with even number of carbon atoms can be formed.

Alkane and CO₂ are liberated at anode while H₂ is liberated at cathode.

(vi) Clemmensen’s reduction

(vii) From compounds containing oxygen Alcohols, aldehydes ketones, carboxylic acids and their derivatives give alkane when treated with hot conc HI and red P in a sealed tube.

(viii) Wolff-Kishnerts reduction

(ix) From carbides

(x) Corey-Bouse synthesis This method can be used to prepare alkanes having odd number of carbon atoms.

Physical Properties of Alkanes

(i) The first four members are colourless gas, next thirteen members are colourless liquids and next higher members are
colourless solids.

(It can be explained on the basis of magnitude of attraction forees.)

(ii). Boiling point of alkanes decreases on branching.

BP ∝ VAF (van der Waals’ forces)

VAF ∝ molecular mass or VAF ∝ SA (surface area)

So boiling point order can be given as

n-octane > iso-octane> 2, 2, 3, 3-tetramethyl butane

(iii). Alkanes with even number of carbon atoms have higher melting points as compared to next higher or lower alkanes with odd number of carbon atoms.

(iv). Alkanes being non-polar in nature, soluble in non-polar solvents but insoluble in polar solvent such as water.

Chemical Properties of Alkanes

(i) Halogenation of alkanes

(a) Chlorination

(b) Bromination of alkanes proceeds in the same way but not so easily.

(c) Iodination

Mechanism of halogenation of alkanes is free radical in nature, i.e., the attacking reagent is a halogen free radical (X^{.). It is a chain reaction.}

(ii) Combustion

Due to the evolution of a large amount of heat during combustion, alkanes are used as fuels.

(iii) Controlled oxidation

(iv) Isomerisation

(v) Aromatisation

Reactions for Methane (CH₄)

(Methane cannot be prepared by Wurtz reaction, Kolbe’s electrolytic process and by reduction of alkenes or alkynes).

Reactions for Ethane(C₂H₆)

Conformations of Alkanes

Alkanes have C-C sigma (σ) bonds and rotation about C-C single bond is allowed. This rotation results in different spatial arrangements of atoms in space which can change into one another, such spatial arrangements are called conformations or conformers or rotamers.

Conformations of ethane

(i) Sawhorse projections

(ii) Newman projections

Intermediate conformation between eclipsed and staggered are known as skew (gauche) conformations.

Eclipsed form is least stable but staggered form is most stable due to greater distance between the bond pairs or lesser torsional strain.

The energy difference between the two extreme forms is of the order of 12.5 kJ mol^-1.

Alkenes

These are unsaturated non-cyclic hydrocarbons which have.sp² -hybridisation with 120° bond angle.

Alkenes are also called olefins [oil.forming] which indicates their high reactive nature.

Alkenes have general formula C_n H_2n, where n = 2,3,4 …

C₂H₄ (ethene), C₃H₆ (propene), etc.

Methods of Preparation of Alkenes

(i) From aIkynes

Physical Properties of Alkenes

Alkene as a class resemble alkanes in physical properties, except in types of isomerism and difference in polar nature.

C₁ to C₃ are gases, the next fourteen are liquids and the higher members are solids.
Alkenes show a regular increase in boiling point with increase in size.

Isomerism in Alkenes

Alkene show both structural isomerism and geometrical isomerism.

Structural isomerism exhibited by alkenes are chain isomerism and position isomerism.

Alkenes also exhibit stereoisomerism as geometrical (cis-trans) isomerism.

Isomerism in Alkenes

Chemical Properties of Alkenes

(i) Addition of halogens

(ii) Addition of hydrogen halides HCI, HBr, HI add up to alkenes to form alkyl halides as per their reactivity order

HI > HBr > HCI

Addition reaction of HBr to unsymmetrical alkenes (Markownikoff’s rule) According to Markownikofrs rule, the negative part of the addendum (adding molecule) gets attached to that carbon atom which possesses lesser number of hydrogen atom.

Anti-Markownikoff addition or peroxide effect or kharash effect In the presence of organic peroxide, addition of only HBr molecule on unsymmetrical alkene takes place contrary to the Markownikoffs rule.

(v) Oxymercuration-demercuration This reaction is an example of hydration of alkene according to Markownikoffs rule.

It is an anti-addition reaction.

It is better than catalytic hydration by dil. H₂SO₄, as it avoids rearrangement.

(vi) Bydroboration oxidation

This reaction involved syn-addition of reagent.

(vii) Oxidation Alkenes decolourise cold dilute aqueous solution of potassium permanganate (Baeyer’s reagent). It is used as a test for unsaturation.

Acidic KMnO₄ or acidic K₂Cr₂O₇ oxidise alkenes to ketones and/or acids depending upon the nature of alkene and the experimental conditions.

(viii) Ozonolysis

(ix) Polymerisation

(x) Reaction with sulphur monochloride

(xii) Diels-Alder reaction

(xiii) Substitution reactions These occur at very high temperature at allylic position

Reactions for Ethene [Ethylene] (C₂ H₄)

Conjugated dienes

Dienes having alternate single (-) and double bonds (=) are called conjugated alkenes. These give Diels, Alder reaction.

Alkynes

These are unsaturated hydrocarbons with general formula C_nH_{2n – 2} e.g., C₂H₂ (ethyne), C₃H₄(propyne)

Alkynes also exhibit electrophilic addition reaction but less reactive than alkenes because the dissociation of x-electron cloud requires more energy.

In alkynes, position of triple bond is determined by ozone (O₃).

H – C = C – H contains 3σ and 2π – bonds and bond length is 120 pm. In acetylene. H – C – C bond angle is 180°.

Methods of Preparation of Alkynes

(i) From calcium carbide

(ii) From vicinal dihalides

(iii) From tetrahalides

Physical Properties of Alkynes

1. The rust two members are gases next eight members (C₅ – C₁₂) are liquids and higher members are solids.
2. They are all colourless and odourless with the exception of acetylene which has slightly garlic odour due to the presence of PH₃ and H₂S as impurities.

3. Alkynes are insoluble in water but soluble in organic solvents like ethers, carbon tetrachloride and benzene.

4. Melting point, boiling point and density increase with increase in molar mass.

Chemical Properties of Alkynes

Alkynes show electrophilic as well as nucleophilic addition reactions.

(i) Acidic character of alkyne

These reactions are not shown by alkenes, alkanes and non-terminal alkynes, hence used for distinction between alkane, alkene and alkyne.

Acetylenic hydrogens are acidic in nature due to 50% a-character in sp-hybridised orbitals.

Acidity of alkynes is lesser than water.

Acidic behaviour order

(ii) Electrophilic addition reactions

The addition product formed depends upon the stability of vinylic cation. Addition on unsymmetrical alkynes takes place according to Markovnikov’s rule.

(a) Addition of dihydrogen

(b) Addition of halogens

(c) Addition of hydrogen halides

(d) Addition of water

(iii) Cyclic polymerisation

(iv) Reaction with AsCl₃ (arsenic trichloride)

(v) Oxidation

Ozonolysis

Higher alkynes give diketones which are further oxidised to carboxylic acid.

(vi) Linear polymerization

Reactions for Acetylene (C₂ H₂)

Benzene

The parent member of the family of aromatic hydrocarbons is benzene (molecular formula: CJls). It has hexagonal ring of six carbon atoms with three double bonds at alternate positions. It is resonance stabilised and the structure may be represented as given ahead.

Methods of Preparation

(i) Cyclic polymerisation of ethyne See alkyne

(ii) Decarboxylation of aromatic acids

Physical Properties of Benzene

Aromatic hydrocarbons are non-polar molecules and are usually colourless liquids or solids with a characteristic aroma.

Aromatic hydrocarbons are immiscible with water but readily miscible with organic solvents.

Aromatic compounds burn with sooty flame.

Chemical Reactions of Benzene

Benzene gives electrophilic substitution reactions.

According to experimental evidences, electrophilic substitution reaction involve following three steps:

(a) Generation of electrophilic
(b) Formation of carbocation intermediate
(c) Removal of proton from the carbocation intermediate.

(i) Nitration

(ii) Halogenation

(iii) Sulphonation

(iv) Friedel-Craft’s alkylation reaction

When Friedel-Craft alkylation is carried out with CH₃CI, the product obtained is C₆H₅CH₃. Incase the alkylation is carried out
with higher alkyl halide, e.g., n-propyl chloride, then the electrophile n-propyl carbocation (CH₃ – CH₂ – ⁺CH₂) which is a
primary carbocation rearranges to form more stable secondary carbocation (iso-propyl carbocation)., and the main product formed will be iso-propyl benzene.

(v) Friedel-Crafts acylation reaction

(vi) With Cl₂ In excess of chlorine, benzene yields hexachlorobenzene [C₆CI₆].

Benzene also undergoes addition reactions e.g.,

Combustion 2C₆H₆ + 15O₂ → 12CO₂ + 6H₂O

Reactions for Benzene

Petroleum

It is a dark coloured oily liquid with offensive odour, found at various depths in many region below the earth’s surface. It is also called rock oil, mineral oil or crude oil It is covered by an atmosphere of a gaseous mixture known as natural gas.

It contains mainly alkanes, cycloalkanes, aromatic hydrocarbons, sulphur, nitrogen and oxygen compounds.

When subjected to fractional distillation, it gives different fractions at different temperatures.

LPG (Liquified Petroleum gas)

It is a mixture of butane and isobutane with a small amount of propane. A strong foul smelling substance, called ethyl mercaptan (C₂H₅SH) is added to LPG cylinders, to help in the detection of gas leakage.

CNG (Compressed Natural Gas)

It consists mainly of methane (95%), which is a relatively unreactive hydrocarbon and makes its nearly complete combustion possible.

Artificial Methods for Manufacturing Petrol

From higher alkanes, petrol or gasoline is obtained by cracking or pyrolysis.

From coal, petrol can be synthesised by following two processes :

(i) Bergius process

(ii) Fischer- Tropsch process

The best catalyst for this process is a mixture of Co, thoria, magnesia and kieselguhr.

The overall yield in this process is slightly higher than Bergius process.

Octane Number

The quality of petrol is expressed in terms of octane number which is defined as the percentage of iso-octane by volume in a mixture of iso-octane and n-heptane which has the same antiknock properties as the fuel under test.

The octane number is 100 for iso-octane (2, 2, 4-trimethylpentane)

Natural gas has octane number 130.

TEL (tetraethyl lead) is used as antiknocking compound.

Octane number is increased by isomerisation, alkylation or aromatisation.

Cetane Number

Quality of diesel oils is measured in terms of cetane number which is defined as the percentage of cetane (hexadecane) by volume in a mixture of cetane and a-methyl naphthalene which has the same ignition property as fuel oil under similar experimental conditions.

It is 100 for cetane and 0 for α – methyl naphthalene.

Methods of Purification of Solids

1. Cystallisation In this process, a saturated solution of impure substance is prepared in hot solvent and heated with animal charcoal which adsorbs the impurities. The solution is filtered and filtrate on cooling deposits crystals of pure compound. Success of the process depends upon the selection of the solvent. The impurities must be least soluble. A process in which crystal formation is initiated by adding crystals of pure substance, is known as seeding.
2. Fractional crystallisation It is based on the different solubilities of different compounds in a solvent. The compound having less solubility crystallises out first on cooling leaving behind others in solution. Sometimes mixture of two solvents, e.g., alcohol and water. chloroform and petroleum ether, give better results.
3. Sublimation Some solids directly convert into vapours when heated without converting into liquid. These are known as sublimate and this process is called sublimation. The substances which sublime can be purified by this method provided the impurities present does not sublime. Camphor, naphthalene and anthracene are purified by sublimation.

Methods of Purification of Liquids

1. Simple distillation The vaporisation of a liquid by heating and subsequent condensation of vapours by cooling is known as distillation. The liquids boiling under ordinary conditions of temperature and pressure without decomposition and containing non-volatile impurities are purified by simple distillation.
2. Fractional distillation It is employed for separating mixture of two or more volatile liquids having boiling points close to each other e.g , acetone (boiling point 60°C) and methanol (boiling point 65°C). Components of petrolium are separated by this method. The vapours of the liquids are passed through the fractionating column which provides greater space for their cooling. The vapours of high boiling substance condense and fall back into distillation flask.
3. Distillation under reduced pressure or vacuum distillation Some liquids decompose when heated to their boiling points e.g., glycerol. Such liquids can be purified by distillation under reduced pressure much below than their boiling points.
4. Steam distillation The liquids insoluble in water, steam volatile in nature, having high molecular weight and high vapour pressure are purified by steam distillation provided the impurities present are not steam volatile. e.g., o-hydroxy acetophenone and jrhydroxy acetophenone are separated by this method.
5. Separating funnel By this, a mixture of two immiscible liquids can be separated.

Chromatographic Method

It was discovered by Tswett (1906).

It is based upon the principle of selective adsorption of various components of a mixture between the two phases: stationary or fixed phase and mobile phase.The various chromatographic
techniques are:

1. Adsorption Chromatography

Stationary phase- Solid or ion exchange resin. Mobile phase-Liquid or gas.

It includes liquid-solid chromatography, gas-solid chromatography or ion exchange chromatography.

2. Partition Chromatography

Fixed phase-liquid supported on inert solid. Mobile phase-liquid or gas.

This process is known as liquid-liquid partition chromatography or liquid-gaspartition chromatography on the basis ofits different phases.

3. Paper Chromatography

The principle of paper chromatography is based on the fact that solutes have the capacity to migrate through filter paper at different rates as a solution is drawn into strip of paper by capillary action.

In paper chromatography, the dissolved substance is applied as a small spot about 2-3 cm from the edge of a strip or square of filter paper and is allowed to dry. This strip is then suspended in a large close container where atmosphere is saturated with the solvent system. The end containing the sample is dipped into the mobile pbase which has already been saturated with the stationary phase. When the solvent front has reached at the other end of the paper, the strip is removed and the zones are located by analytical methods.

(The ratio of the distance travelled by a component to the distance travelled by the solvent front is characteristic of each component and is known as the R_f value.

R_f = (distance in em from starting line to the centre of zone/distance in cm from starting line to the solvent front)

4. Column Chromatography

It is an example of adsorption chromatography. Adsorbents used are alumina, silica gel, cellulose powder, animal charcoal, keiselguhr etc.

Liquid solvents used are benzene. petroleum ether, alcohol etc.

When the solvent is poured over the mixture present at the top of a column packed with adsorbent, the components are separated into, number of layers called zones, bands or chromatograms due to preferential adsorption.

Elution The continuous pouring of solvent from the top of the column is known as elution or running of column. Solvent is known as eluant. The most weakly adsorbed component is eluted first by least polar solvent while more strongly adsorbed component is eluted later by highly polar solvents.

Chemical Methods of Purification

The substance to be purified is treated with a suitable chemical reagent to form a stable derivative. It is then separated by suitable method and decomposed to get the pure compounds.

Examples

1. Mixture of amines (1°, 2° and 3°) is separated by Hinsberg’s method.
2. Acetic acid from pyroligneous acid is separated by forming calcium salt.
3. Acids are separated by froming sodium derivatives with NaHCO₃·
4. Absolute alcohol is obtained from rectified spirit by quick lime process and azeotropic distillation.

1. Azeotropic Distillation

Azeotropes are constant boiling mixtures which distil off without any change in composition at a fixed temperature. Therefore, components of an azeotropic mixture cannot be separated by fractional distillation. A very common example of azeotropic mixture is rectified spirit which contains 95.87% ethyl alcohol and 4.23% water by weight which boils at 351.1 K.

(Such mixtures are separated by adding another component which generate a new lower boiling azeotrope that is heterogeneous (i.e., producing two immiscible liquid phases). e.g., C₆H₆ is added to H₂O and ethyl alcohol azeotrope to separate them.

2. Differential Extraction

This method is used to separate an organic compound present in aqueous solution which is more soluble in other solvent than in water.

Qualitative Analysis of Organic Compounds

1. Detection of Carbon and Hydrogen

This is done by heating the given organic compound with dry cupric oxide in a hard glass test tube when carbon present is oxidised to carbon dioxide and hydrogen is oxidised to water.

Carbon dioxide turns lime water milky.

Water condenses on the cooler parts of the test tube and turns anbydrous copper sulpbate blue.

Lassaigne’s Test

The organic compound is fused with a small piece of Na metal. When
element (N, 8, X) of the organic compound combine to give NaCN,Na₂S or NaX; the red hot tube is plunged in distilled water, boiled and filtered. The filtrate is called Lassaigne’s extract or sodium extract. The Lassaigne’s extract is usually alkaline. If not, it is made alkaline by adding a few drops of a dilute solution of sodium hydroxide.

The purpose of fusing the organic compounds with sodium metal is to convert halogens, N, S, P etc., present in the organic compound to their corresponding soluble sodium salts (ionic compounds).

Na + C + N → NaCN
2Na + S → N₂S (where, X = Cl, Br, I)
Na + X → NaX

1. Detecton of Nitrogen

To a part of this alkaline solution is added a few drops of a freshly prepared solution of ferrous sulphate, because a dilute solution of FeSO₄ after a long time oxidise to basic ferric sulphate which is useless for analysis. The contents are warmed a little, cooled and then acidified with dil. H₂SO₄. Appearance of a green or Prussian blue colouration indicates the presence of nitrogen.

If S is also present alongwith N, a red colour in place of Prussian blue in the test of nitrogen appears, due to the formation of Fe(CNS)₃.

Hydrazine does not give Lassaigne’s test for nitrogen since it does not contain carbon. In order to test the presence of N in such compounds, during fusion with Na, some charcoal or preferably strch (which contains C but not N, S, halogens etc.) is added. Under these conditions. C of starch or charcoal combines with N of the compound to form NaCN which will now give a positive test for nitrogen.

Lassaigne’s test is not shown by diazonium salts because diazonium salts usually lose N₂ on heating much before they have a chance to I react with fused sodium metal.

2. Detection of Sulphur

(i) Sodium fusion extract is acidified with acetic acid and acetate is added to presence of sulphur.

(ii) On treating sodium fusion extract with sodium nitroprusside, apperance of a violet colour further indicates the presence of sulphur.

3. Detection of Halogens

The sodium fusion extract is acidified with nitric acid and then treated with silver nitrate.

X^– + Ag⁺ → AgX

X represents a halogen -Cl Br, or I.

AgCl white ppt, AgBr-dull yellow ppt, AgI-bright yellow ppt.

Note Beilstein test is also a test for halogen but it is not a confirmatory test.

Detection of Phosphorus

The compound is heated with an oxidising agent (sodium peroxide). By this the phosphorus present in the compound is oxidised to phosphate.

The solution is boiled with nitric acid and then treated with ammonium molybdate. A yellow colouration or precipitate indicates the presence of phosphorus.

Detection of Oxygen

There is no direct method to detect oxygen in compounds. It is present in the form of functional groups such as -OH, -COOH. -NO₂ etc.

Qauantitative Estimation of Elements

1. Estimation of Carbon and Hydrogen (Liebig’s Method) When a known mass of organic compound is strongly heated with dry euO, C and H present are quantitatively oxidised to CO₂and H₂O respectively,

By knowing the amount of CO₂ and H₂O from known weight of organic compound, the percentage of carbon and hydrogen can be computed,

The water is absorbed in anhydrous CaCl₂.

The carbon dioxide is absorbed in concentrated solution of KOH.

.

On heating with CuO, elements other than C and H are also modified as follows:

When organic compound contains nitrogen. the oxides of nitrogen

(NO,N₂O etc.) are absorbed by caustic potash. These are removed by the use of bright copper gauge.

4Cu + 2NO₂ → 4CuO + N₂

Cu + N₂O → CuO + N₂

Nitrogen is not absorbed by KOH solution.

When organic compound contains halogens, they are removed by using silver gauge by forming non-volatile silver halide.

When sulphur is present, it is removed by forming lead sulphate by using fused lead chromate and halogens form lead halides.

Estimation of Nitrogen

(i) Duma’s method This method is used for nitrogenous compounds. Though tedious but it is better than Kjeldahl’s method.

In this method, the nitrogenous compound is heated strongly with ceo in the atmosphere of CO₂and the mixture obtained is passed over a roll of heated bright Cu gauze. The oxides of nitrogen again reduce to N₂. The resultant mixture is passed in KOH. All gases except N₂ are fairly absorbed. Nitrogen is collected over KOH and its volume at NTP is measured.

Organic compound + conc. H₂SO₄ + (small amount of K₂SO₄ and its volume at NTP is measured.

(ii) Kjeldahl’s method Organic compound + conc. H₂SO₄ + (small amount of K₂SO₄ and

Ammonia is passed through H₂SO₄ or HCl of known volume and normality. The volume of acid neutralised by NH₃ is calculated by neutralising the acid left by NaOH solution.

Percentage of nitrogen = (1.4 x N x V/wt. of organic compound)

N = normality of acid
V = volume of acid in mL neutralised by ammonia.

(In practice, K₂SO₄ is added to raise the boiling point of H₂SO₄ and CuSO₄ is added to catalyse the reaction).

Kjeldahl’s method is not reliable as results obtained are generally low. It cannot be applied to compounds containing nitrogen directly linked to oxygen or nitrogen such as nitro, nitroso, azo and nitrogen present in ring as in pyridine.

Estimation of Halogen (Carius Method)

Organic compound + fuming HNO₃ + AgNOa → AgX

It is estimated gravimetrically.

Percentage of halogen = (wt. of halogen atom x wt. of AgX x 100/mol. wt. of AgX x wt. of organic compound)

Estimation of Sulphur

Estimation of Phosphorus

Organic compound + Fuming nitric acid → H₃PO₄

Now a day CHN elemental analyser is used to estimate the C, H and N in the organic compound.

Determination of Empirical Formula

Empirical formula expresses the relative number of atoms present in the molecule. It is calculated from percentage composition of the compound.

Determination of Molecular Formula

Molecular formula = (Empirical formula)_n

n = (molecular weight/empirical formula weight)

Classification of Environment

1.Atmosphere

Atmosphere is Ii gaseous mixture of air that surrounds the earth. Its different layers are as

(1) Troposphere It is the lowest region of the atmosphere extending from earth’s surface to the lower boundary of the stratosphere. It contains water vapours and is greatly affected by air pollution.

(ii) Stratosphere The layer of the earth’s atmosphere above the troposphere and below the mesosphere, is called stratosphere. Ozone layer to; present in this region.

(iii) Mesosphere It is the region of the earth’s atmosphere above the stratosphere and below the thermosphere. It is the coldest region (temperature – 2 to 92°C) of atmosphere.

(iv) Thermosphere The upper region of the atmosphere above the mesosphere is called thermosphere It is the hottest region (temperature up to 1200°C).

(v) Exosphere It is the uppermost region of atmosphere. It contains atomic and ionic O₂, H₂ and He.

2. Hydrosphere

It is the aqueous envelop of the earth e.g., oceans. lakes etc.

3. Lithosphere

The solid rocky portion of the earth constitute the lithosphere.

4. Biosphere

The biological envelop which supports the life is called biosphere. e.g., animal, human beings.

Environmental Pollution

It may be described as contamination of environment with harmful wastes mainly arising from certain human activities. These activities release materials which pollute atmosphere, water and soil.

Types of Pollutions

(i) Natural pollution This type of pollution is caused by the natural sources e.g., volcanic eruptions. release of methane by paddy fields and cattles, forest fires etc.

(ii) Man-made pollution This type of pollution is resulting from human activities like burning of the fuels, deforestation, industrial effluents, pesticides etc.

Pollutants

Any substance produced either by a natural source or by human activity which causes adverse effect on the environment is called pollutant.

Pollutants can be of the following types depending upon the following factors:

Classification on the Basis of Their Degradation

(i) Biodegradable pollutants Pollutants capable of being degraded by biological or microbial actions are called biodegradable pollutants, e.g., domestic sewage.

(ii) Non-biodegradable pollutants The substances which are normally not acted upon by microbes are called non-biodegradable pollutants. These undergo biological magnification.

They can further be of two types

(i) Wastes, e.g., glass, plastic, phenols

(ii) Poisons, e.g., radioactive substances, Hg salts, pesticides. heavy metals.

Classification on the Basis of Their Occurrence in Nature

(i) Primary pollutants These are present in same form in which these are added by man e.g., DDT. pesticides. fertilizers etc.

(ii) Secondary pollutants These occur in different forms and are formed by the reaction between the primary pollutants in the presence of sunlight e.g., HNO₃, H₂SO₄ PAN, ozone etc.

Classification on the Basis of Their Existence in Nature

(i) Quantitative pollutants These are naturally present in nature and also added by man. These become pollutants when their concentration reaches beyond a threshold value in the environment, e.g., CO2, nitrogen oxide etc.

(ii) Qualitative pollutants These arc not present in the nature but are added by nature only due to human activities. e.g., pesticides. fungicides. herbicides etc.

Tropospheric Pollution

It is caused by gaseous pollutants and particulate matter.

Gaseous air pollutants Oxides of sulphur (SO_x), oxides of nitrogen (NO_x), oxides of carbon (CO, CO₂), hydrogen sulphide (H₂S), hydrocarbons etc.

Particulate pollutants Dust, fumes. mist, smoke etc.

Air Pollution

Air pollution occurs when the concentration of a normal component of the air or a new chemical substance added or formed in air, build up to
undesirable proportions causing harm to humans, animals, vegetation and materials. The chemical substances and particles causing pollution are called air pollutants.

Air Pollutants

The major air pollutants are

(i) Carbon monoxide (CO) It is produced by incomplete combustion of gasoline in motor vehicles, wood. coal, incineration and forest fires. It induces headache, visual difficulty, coma or death. It blocks the normal transport of oxygen from the lungs to other parts of the body, by combining with haemoglobin of the blood. (Its affinity towards haemoglobin is about 200 times more than the oxygen.)

(ii) Sulphur dioxide (SO₂) It is produced by petrol combustion, coal combustion, petrol refining and smelting operation.

It obstruct the movement of air in and out of lungs. It is particularly poisonous to trees causing chlorosis and dwarfing. In the
presence of air. it is oxidised to SO₃ which is also an irritant.

2SO₂ + O₂ (air) → 2SO₃

Taj Mahal is reported to be affected by SO₂ and other pollutants released by oil refinery of Mathura.

(iii) Oxides of nitrogen NO₂ and NO are obtained by combustion of coal, gasoline. natural gas. petroleum refining, chemical industries and tobacco smoke. In upper atmosphere. these are emitted by high flying jets and rockets.

Breathing NO₂ causes chlorosis to plants and chronic lung conditions leading to death in human beings . These oxides destroy ozone layer.

(iv) Smoke, dust These are obtained in cement works, iron and steel works. gas works, power generating stations. Coal miners
suffer from black lung disease and textile workers suffer from white lung disease.

(v) Ammonia It is produced by fertilizer works.

(vi) Mercaptans These are obtained from oil refineries. coke ovens etc.

(vii) Zn and Cd These are obtained from zinc industries.

(viii) Freon (or CFC’8) Their source is refrigerator.

Smog

It is a mixture of smoke (composed of tiny particles of carbon, ash and oil etc from coal combustion) and fog in suspended droplet form. It is of
two types:

1. London smog or classical smog

It IS coal smoke plus fog The fog part is mainly SO₂ and SO₃. It has sulphuric acid aerosol. It causes bronchial irritation and acid rain. It is reducing in nature and occurs in cool humid climate.

2. Photochemical smog or Los Angeles smog

The oxidised hydrocarbons and ozone In a warm. dry and sunny climate cause photochemical smog. Its brown colour is due to the presence of NO₂.

The nitrogen dioxide by absorbing sunlight in blue and UV region decomposes into nitric oxide and atomic oxygen followed by a series of the other reactions producing O₃, formaldehyde, acrolein and peroxyacetyl nitrates.

Hydrocarbons + O₂, NO₂ NO, O, O₃ Peroxides, formaldehyde, peroxyacetyl nitrate (PAN), acrolein etc.

It is oxidising in nature and causes irritation to eyes, lungs, nose, asthmatic attack and damage to plants.

Green House Effect and Global Warming

The phenomenon in which atmosphere of earth traps the heat coming from the sun and prevents it from escaping into the outer space is called green house effect. Certain gases, called green house gases [carbon dioxide, methane, ozone, chlorofluorocarbon compounds (CFCs) and water vapour] in the atmosphere absorb the heat given by earth and radiate back it to the surface of the earth. Thus, warming of the earth led to the warming of air due to green house gases. which is called global warming.

Consequences of Green House Effect (or Global Warming)

1. The green house gases are useful in keeping the earth warm with an average temperature of about 15° to 20°C.

2. There may be less rainfall in this temperature zone and more rainfall in the dried areas of the world.

3. Increase in the concentration of CO₂ in the atmosphere leads to increase in the temperature of the earth’s surface. As a result evaporation of surface water will increase which further help in the rise of temperature and results in the melting of glaciers and polar ice caps and hence, level of sea water may rise.

ACid Rain

The pH of normal rain water is 5.6 due to the dissolution of CO₂ from atmosphere.

when the pH of rain water drops below 5ppm, it is called acid rain (by Robert Augus.) Oxides of N ans S are responsible for making rain
water acidic, Much of the NO_x and SO_x entering in the atmosphere are converted into HNO₃ and H2SO₄ respectively. The detailed photochemical reactions occurring in the atmosphere are given as

HNO₃ is removed as a precipitate or as particulate nitrates after reaction with bases (like NH₃, particulate lime etc).

The presence of hydrocarbons and NO_x step up the oxidation rate of the reaction. Soot particles are also known to be strongly involved in catalysing the oxidation of SO₂

Acid rain causes extensive damage to building and sculptural materials of marble, limestone, slate. mortal’ etc

Stratospheric Pollution (Depletion of Ozone Layer)

Ozone is a light bluish gas and absorbs UV radiations of the sun which are harmful to living beings, But nowadays ozone layer is being depleted by CFCs (chlorofluorocarbons).

UV radiations cause the chlorofluorocarbons to dissociate to fOl1D highly reactive chlorine free radical which reacts with ozone to form chlorine monoxide.

CI* (free radical) can react WIth more O₃.

Ozone hole is formed over Antarctica. and some parts of non – polar regions also.

In other parts of stratosphere NO₂, CH₄ react with CIO* and cI* respectively and act as natural sink for CIO* and CI*

These reactions consume CI* and CIO* hindrance to ozone depletion.

[In Antarctica, during winters, special types of clouds, called polar stratospheric clouds (PSCs)are formed. These clouds are of two types

Type I Clouds They contain some solidified nitric acid trihydrate (HNO₃ * 3H₂O) formed at about -77°C.

Type II Clouds They contain some ice formed at about – 85°C. These clouds play important role in ozone depletion by hydrolysing
chlorine nitra.

Hypochlorous acid and CI₂ are formed which are reconverted into reactive chlorine atoms with the help of sunlight which causes ozone depletion.]

Polar vortex During winters, when polar stratospheric clouds are formed over Antarctica. stable wind patterns in the stratosphere encircle the continent which is called polar vortex. It is tight whirlpool of winds which is so rigid that air within it is isolated from the sun and forms the warmer air of temperate region to fill up ozone hole.

Consequences of Depletion of Ozone Layer

(a) Loss of sight The UV radiation damage the cornea and lens of the eyes.

(b) Effect on immune system The UV radiations are also likely to suppress immune system.

(c) Skin cancer This type of radiation is known to be cancer causing agent.

Water Pollution

The contamination of water by foreign substances which would COnstitute a health hazard and make It unfit for all purposes (domestic, industrial or agriculture etc) is known as water pollution. The polluted Water may have foul odour. bad taste, unpleasant colour etc.

Maximum prescribed concentration of some metals in drinking water is as

Sources of Water Pollution

(i) Domestic sewage Discharge from kitchens, baths, etc.

(ii) Industrial water Wastes from manufacturing processes which includes acids, alkalies, pesticides, insecticides, metals. fungicides etc.

(iii) Oil From oil spills or washings of automobiles.

(iv) Atomic explosion Processing of radioactive materials.

(v) Suspended particles (organic or inorganic) Viruses, bacteria, algae, protozoa etc.

(vi) Wastes from fertilizer Industries such as phosphates, nitrates, ammonia etc.

(vii) Clay Ores, minerals, fine particles of soil.

Effects of Impurities in Water

(a) Fluorides Mottling of teeth enamel, above 1 mg/L fluoride causes fluorosis.

(b) Sulphates Sulphates of Na, K, Mg cause diarrhoea.

(c) Lead It damages kidney, liver, brain and central nervous system.

(d) Cadmium and mercury They causes kidney damage.

(e) Zn It causes dizziness and diarrhoea. .

(f) Arsenic It can cause cramps and paralysis.

(g) Phosphates from fertilizers They promote algae growth and reduce dissolved oxygen concentration of water. This process is known as eutrophication.

Aerobic and Anaerobic Oxidation

The oxidation of organic compounds present in sewage in the presence of good amount of dissolved or free oxygen (approx, 8.5 mlJL) by aerobic bacteria is called aerobic oxidation. When dissolved or free oxygen is below a certain value, the sewage is called stale.

Anaerobic bacteria bring out putrefaction by producing H₂S, NH₃, CH₄, (NH₄)₂S etc. This type of oxidation is called anaerobic oxidation.

The optimum value of dissolved oxygen for good quality of water is 4·6 ppm (4-6 mg/L). The lower the concentration of dissolved oxygen, the more polluted is the water.

Biological oxygen demand (BOD) It is defined as the amount of free oxygen required for biological oxidation of the organic matter under aerobic conditions at 20°C for a period of five days. Its unit is mg/L or ppm.

An average sewage has BOD of 100 to 150 mg/L.

Chemical oxygen demand (COD) It is the measure of all types of oxidisable impurities (biologically oxidisable and biologically inert organic matter such as cellulose) present in the sewage. COD values are higher than BOD values.

Control of Water Pollution

(i) Recycling of waste water

(ii) Use of chemicals Lead poisoning can be cured by giving the patient an aqueous solution of calcium complex of EDTA. Lead ions displace calcium in the EDTA complex to form chelated lead and Ca²⁺. The soluble lead chelate is excreted with the urine.

Ca – EDTA + Pb²⁺ → Pb – EDTA + Ca²⁺

(iii) Special techniques such as adsorption, ion exchangers, reverse osmosis, electrodialysis etc.

(iv) Waste water reclamation

Sewage Treatment

It involves the following steps

(i) Preliminary process Passing sewage through screens to remove large suspended matter and then through mesh screens to remove solids, gravels, silt etc.

(ii) Settling process (sedimentation) The residual water when allowed to stand in tanks, the oils and grease, float on the surface and skimmed off and solids settle down. The colloidal material is removed by adding alum, ferrous sulphate etc. Primary sludge can be separated.

(iii) Secondary treatment or biological treatment It is aerobic chemical oxidation or aeration which converts carbon of the organic matter to CO₂, nitrogen into NHJ and finally into nitrite and nitrates, dissolved bases form salts such as NH₄O₂, NH₄NO₃ and Ca(NO₃)₂ etc., and secondary sludge is obtained.

(iv) Tertiary treatment It is treatment of waste water with time for removal of phosphate which IS then coagulated by adding alum and ferric chloride and removed by filtration.

Water is disinfected by adding chlorine.

Secondary sludge forms a good fertilizer for soil as it contains nitrogen and phosphorus compounds.

Soil or Land Pollution

The addition of substances in an indefinite proportion changes the productivity of the soil. This is known as soil or land pollution.

Sources of Soil Pollution

(i) Agricultural pollutants e.g., chemicals like pesticides, fertilizers, bactericides, fumigants. insecticides, herbicides, fungicides.

(ii) Domestic refuge and industrial wastes.

(iii) Radioactive wastes from research centres, and hospitals.

(iv) Soil conditioners containing toxic metals like Hg, Pb, As. Cd etc.

(v) Farm wastes from poultries, dairies and piggery farms.

Control of Soil Pollution

(i) Use of manures Manures prepared from animal dung is much better than the commonly used fertilizers.

(ii) Use of bio- fertilizers These are the organisms which are inoculated in order to bring about nutrient enrichment of the soil. e.g., nitrogen fixing bacteria and blue-green algae.

(iii) Proper sewerage system A proper sewerage system must be employed and sewage recycling plants must be installed.

(iv) Salvage and recycling Rag pickers remove a large number of waste articles such as paper, polythene, card board. rags. empty bottles and metallic articles. These are subjected to recycling and this helps in checking soil pollution.

Radioactive Pollution

Cosmic rays that reach the parth from outer space and terrestrial radiation from radioactive elements are natural radiations. This natural or background radiation is not a health hazard due to its low concentration.

Man made sources of radiations include mining; and refining of plutonium and thorium, atomic reactors and nuclear fuel. These are produced during preparation of radio-isotopes. These are of two types: electromagnetic (radio waves UV, IR, α-rays) and particulate.

Other Sources of Radioactive Pollution

(i) Atomic explosions Atomic explosions produce radioactive particles which are thrown high up into the air as huge clouds.

The process releases large amount of energy as heat. Due to atomic -explosion nuclear fallout. These radioactive elements may reach the human beings through food chain.

(ii) Radioactive wastes Wastes from atomic power plants come in the form of spent fuels of uranium, and plutonium. People
working in such power plants, nuclear reactors, fuel processors etc., are vulnerable to their exposure.

(iii) Radio isotopes Many radioactive isotopes like C¹⁴, I¹²⁵, p³² and their compounds are used in scientific researches. The waste water of these research centres contains the radioactive elements which may reach the human beings through water and food chains.

Effects of Radiations

1. Strontium-90 accumulates in the bones to cause bone cancer and tissue degeneration in number of organs.

2. 1-131 damages WBCs, bone marrow, lymph nodes and causes skin cancer, sterility and defective eye sight.

3. These may cause ionisation of various body fluids, chromosomal aberrations and gene mutations.

4. Radioactive iodine may also cause cancer of thyroid glands.

5. Cesium-137 brings about nervous, muscular and genetic change.

6. Uranium causes skin cancers and tumours in the miners.

7. Radon-222 causes leukemia, brain tumours and kidney cancers.

Bhopal Gas Tragedy

In Dec. 2, 1984 a dense cloud of methyl isocyanate gas (Mlq leaked from a storage tank of the Union Carbide ltd plant in Bhopal. It caused a great loss of life to people and animals. Methyl Isocyanate was prepared by the reaction of methyl amine with phosgene and stored in abundance

Green Chemistry-An Alternative Tool for Reducing Pollution

Green chemistry may be called chemistry involved in the design, development, and implementation of chemical products and processes to reduce or eliminate the use and generation of substances hazardous to human health and the environment.

Thus, the goal of green chemistry is ‘to promote the development of products and processes that reduce or eliminate the use or generation of toxic substances associated with the design, manufacture, and use of hazardous chemicals. Some important principles and method of green chemistry are

1. It is better to prevent waste than to treat or clean up waste after it is formed.

2. Synthetic methods should be designed to maximize the incorporation of all materials used in the process into the final product. .

3. Whenever possible, synthetic methodologies should be designed to use and generate substances that possess little or no toxicity to human health and the environment.

4. Chemical products should be designed to preserve efficiency of function while reducing toxicity.

5. The use of auxiliary substance (e.g., solvents, separation agents etc.) should be avoided as far as possible.

6. Energy requirements should be recognised for their environmental and economic impacts and should be minimized.

7. Synthetic methods should be conducted at ambient temperature and pressure.

Group 13

It is also called boron family. It includes B, Al, Ga, In, Tl. AI is the most abundant metal and third most abundant element in the earth’s crust.

General Physical Properties of Group 13 Elements

(i) Electronic configuration Their valence shell electronic configuration is ns²np¹

(ii) Atomic radii and ionic radii Group 13 elements have smaller size than those of alkaline earth metals due to greater effective nuclear charge, Z_eff’

Atomic radii increases on going down the group with an anomaly at gallium (Ga). Unexpected decrease in the atomic size of Ga is due to the presence of electrons in d-orbitals which do not screen the attraction of nucleus effectively.

The ionic radii regularly increases from B³⁺ to TI³⁺.

(iii) Density It increases regularly on moving down the group from B to Tl.

(iv) Melting and boiling points Melting point and boiling point of group 13 elements are much higher than those of group 2 elements. The melting point decreases from B to Ga and then increases, due to structural changes in the elements.

Boron has a very high melting point because of its three dimensional structure in which B atoms are held together by strong covalent bonds.

Low melting point of Ga is due to the fact that it consists of Ga₂ molecules, and Ga remains liquid upto 2276 K. Hence, it is used in
high temperature thermometer.

(v) Ionisation enthalpy (IE) The first ionisation enthalpy values of group 13 elements are lower than the corresponding alkaline earth metals, due to the fact that removal of electron is easy. [ns² npl configuration] .

On moving down the group, IE decreases from B to Al, but the next element Ga has slightly higher ionisation enthalpy than A1 due to the poor shielding of intervening d-electrons. It again decreases in In and then increases in the last element Tl

(vi) Oxidation states B and Al show an oxidation state of +3 only while Ga, In and TJ exhibit oxidation states of both +1 and +3.

As we move down in the group 13. due to inert pair effect, the tendency to exhibit +3 oxidation state decreases and the tendency to attain +1 oxidation state increases.

Stability of +1 oxidation state follows the order Ga < In < Tl.

Inert pair effect is reluctance of the s-electrons of the valence shell to take part in bonding. It occurs due to poor shielding of the ns² – electrons by the intervening d and f – electrons. It increases down the group and thus, the lower elements of the group exhibit lower oxidation states.

(vii) Electropositive (metallic) character These elements are less electropositive than the alkaline earth metals due to their smaller size and higher ionisation enthalpies.

On moving down the group, the electropositive character first increases from B to Al and then decreases from Ga to TI, due to the presence of d and I-orbitals which causes poor shielding.

(viii) Reducing character It decreases down the group from AI to Tl because of the increase in electrode potential value for M³⁺ / M.

Therefore, it follows the order

AI> Ga > In > Tl

(ix) Complex formation Due to their smaller size and greater charge, these elements have greater tendency to form complexes than the s-block elements.

(x) Nature of compounds The tendency of the formation of ionic compounds increases from B to Tl. Boron forms only covalent compounds whereas AI can form both covalent as well as ionic compounds. Gallium forms mainly ionic compounds, although anhydrous Ga CI₃ is covalent.

Chemical Properties of 13 Group Elements

(i) Action of air Crystalline boron is unreactive whereas amorphous boron is reactive. It reacts with air at 700°C as follows

4B + 3O₂ → 2B₂O₃

2B + N₂ → 2BN

AI is stable in air due to the formation of protective oxide film.

4Al + 3O₂ → 2Al₂O₃

Thallium is more reactive than Ga and In due to the formation of unipositive ion, TI⁺.

4Tl + O₂ → 2Tl₂0

(ii) Reaction with nitrogen

(iii) Action of water Both B and AI do: not react with water but amalgamated aluminium react with H₂O evolving H₂.

2Al(Hg) + 6H₂O )→ 2AI(OH)₃ + 3H₂ + 2Hg

Ga and In do not react with pure cold or hot water but Tl forms an oxide layer on the surface.

(iv) Reaction with alkalies Boron dissolves in alkalies and gives sodium borates.

Aluminium also reacts with alkali and liberates hydrogen.

(v) Reaction with carbon

Aluminium carbide is ionic and forms methane with water.

(vi) Hydrides Elements of group 13 do not combine directly with H₂ to form hydrides, therefore their hydrides have been prepared by indirect methods, e.g

Boron forms a number of hydrides, they are known as boranes. Boranes catch fire in the presence of oxygen.

B₂H₆ + 3O₂ → B₂O₃ + 3H₂O; &Delat;_cH° = – 1976 kJ mol^-l

Boranes are hydrolysed by water.

B₂H₆ + 6H₂O → 2H₃BO₃ + 6H₂

Boranes are stable but the stability of hydrides of AI, Ga, In, and Tl decreases on moving down the group because the strength of the M-H bond decreases.

Structure of diborane BH₃ does not exist as such, but exists as a dimer, i.e; B₂H₆(diborane].

In the above structure, B atoms are in Sp³ – hybrid state. There are six B-H bonds out of which four B-H bonds are normal bonds present in the same plane While rest two B-H bonds behave as bridge bonds, ie; 3c – 2e (three centre-two electrons, also known as banana bond) and present above and below the plane of the molecules which do not I have sufficient number of electrons to form covalent bonds.

Aluminium (AI) forms a polymeric hydride of general formula (AIH₃)_x which decomposes into its elements on heating.

(vii) Oxides Except. TI all the elements of group 13 form oxides or general formula M₂O₃ on heating with oxygen.

TI forms thallium (l) oxide. Tl₂O which IS more stable than thallium (III) oxide TI₂O₃ due to inert pair effect.

(viii) Nature of oxides and hydroxides B(OH)₃ or H₃BO₃ is soluble in water, while other hydroxides are insoluble in water.

On moving down the group. there is a change from acidic to amphoteric and then to basic character of oxides and hydroxides or group 13 elements.

(ix) Halides All the elements of boron family (except Tl) form trihalides of type MX₃.

All the boron trihalides [BX₃) and aluminium trihalides AlX₃ (except AIF₃ which is ionic) are covalent compounds. AlX₃ exists as dimer while BX₃ is monomer because boron atom is too small to coordinate with four large halide ions. The energy released during the formation of the bridge structure is not sufficient for the cleavage of the typical pπ – pπ bond in BF₃.

BF₃ is a colourless gas, BCl₃ and BBr₃ are colourless fuming liquids and BI₃ is a white solid at room temperature.

Trihalides of group 13 elements behave as Lewis acids because of their strong tendency to accept a pair of electrons. The relative strength of Lewis acids of boron trihalides is

BF₃ < BCI₃, < BBr₃, < BI₃.

This is due to pπ – pπ backbonding in BF₃ which makes it less electron deficient.

The halides of group 13 elements behave as Lewis acids and the acidic character is

BX₃ > AIX₃ > GaX₃ > InX₃ (where, X = Cl, Br or 1)

TICI₃ decomposes to TICl and C1₂ and hence acts as an oxidising agent.

Anomalous Behaviour of Boron

Boron shows anomalous behaviour with the other members of the group, due to the following reasons:

(i) Smallest size in the group.
(ii) High ionisation energy.
(iii) Highest electronegativity in the group.
(iv) Absence of vacant d-orbital.

A few points of difference are

1. It is a non-metal while other members of the group are metallic.
2. It shows allotropy while other members do not.
3. It has the highest melting point and boiling point in group 13.
4. It forms only covalent compounds while other members form both ionic and covalent compounds.
5. The halides of boron exist as monomers while AlCI:! exists as a dimer.
6, The oxides and hydroxides of boron are weakly acidic while those of aluminium arc amphoteric and those of other elements are basic.
7. It can be oxidised by concentrated HNO₃ while aluminium becomes passive due to the formation of oxide layer on the surface.

Diagonal Relationship between Boron and Silicon

Boron exhibit resemblance with its diagonal element silicon of group 14.

1. Both Band Si are non-metals.
2. Both arc semi-conductors.
3. Both Band Si form covalent hydrides, i.e.. boranes and silanes respectively.
4. Both form covalent, and volatile halides which fume in moist air due to release of HCI gas.

BCI₃ + 3H₂O → H₃ BO₃ + 3HCl i

SiCl₄ + 4H₂O → Si(OH)₄ + 4HCl

5. Both form solid oxides which get dissolve in alkalies forming borates and silicates respectively. ..
6. Both react with electropositive metals and give binary compounds, which yield mixture of boranes and silanes on hydrolysis.

Boron and Its Compounds

Occurrence

It does not occur in free state. Its important minerals are

(i) Borax (or Tineal), Na₂B₄O₇ * 1OH₂O
(ii) Kernite, Na₂B₄O₇ * 4H₂O
(iii) Orthoboric acid, H₃BO₃

Isolation

Elemental boron is obtained by following methods :

(i) By reduction of boric oxide with highly electropositive metals like K, Mg, AI, Na etc, in the absence of air.

(ii) By the reaction of boron halides with hydrogen,

Uses of Boron

(i) As a semi-conductor.
(ii) Boron steel rods are used to control the nuclear reactions.

₅B¹⁰ + ₀n¹ → ₅B¹¹

1. Borax or Sodium Tetraborate Decahydrate [Na₂B₄O₇ * 1OH₂O]

Preparation

It occurs naturally as tineal in dried up lakes. It is obtained by boiling of mineral colemanite with a solution of Na₂Co₃.

NaBO₂ can be removed by passing CO₂ through it.

4NaBO₂ + CO₂ → Na₂CO₃ + Na₂B_<4O₇

Properties

1. Its aqueous solution is basic in nature.

Na₂B₄O₇ + 7H₂O → 2NaOH + 4H₃BO₃

2. On heating with ethyl alcohol and cone.H₂SO₄it gives volatile vapours of triethyl borate which burn with a green flame.

3. Action of heat

Borax bead is used for the detection of coloured basic radicals under the name borax bead test e.g.,

2. Boric Acid or Orthoboric Acid [H₃BO₃ or B(OH)₃]

Preparation

By treating borax with dil. HCl or dil. H₂SO₄.

Na₂B₄O₇ + 2HCl + 5H₂O → 2NaCI + 4H₃BO₃

Properties

1. It is a weak monobasic acid (Lewis acid).

H₃BO₃ + 2H₂O → [B(OH)₄]^– + H₃O⁺

2. With C₂H₅OH and cone H₂SO₄, it gives triethyl borate.

Uses

It is used as an antiseptic and eye lotion under the name ‘boric lotion’, and as a food preservative.

3. Borazine or Borazole, [B₃N₃H₆]

It is a colourless liquid having a six membered ring of alternating B and N atoms. It is also called ‘inorganic benzene’. It is prepared by B₂H₆ as follows

The π electrons in borazine are only partially delocalised. It is more reactive than benzene

Compounds of Aluminium

1.Anhydrous Aluminium Chloride [AICI₃ or A1₂C1₆]

Preparation

It can not be prepared by heating AICI₃. 6H₂O.

It can be prepared

(i) By passing dry chlorine or HCl gas over heated Al.

(ii) By heating a mixture of alumina and carbon in a current of dry chlorine.

Properties

1. AlC1₃ fumes in moist air due to hydrolysis.

AlC1₃ + 3H₂O → Al(OH)₃ + 3HCI

2. It behaves as Lewis acid.

Uses

It is used as a catalyst in Friedel-Craft reaction and as a mordant dye.

2.Aluminium Oxide or Alumina [AI₂O₃]

It is the most stable compound of aluminium and occurs in nature as colourless corundum and several coloured oxides, (it present in combination with different metal oxides) like ruby (red), topaz (yellow), sapphire (blue), and emerald (green), which are used as precious stones (gems).

Alum

The term alum is given to double sulphates of the type X₂SO₄ * Y₂(SO₄)₃ * 24H₂O where, X represents a monovalent cation such as Na⁺, K⁺ and NH⁺₄, while Y is a trivalent cation such a Al³,Cr³⁺, Fe³⁺ and Co³⁺(Li⁺ does not form alum).

Some important alums are

(i) Potash alum K₂SO₄ * Al₂(SO₄)₃ * 24H₂O

(ii) Sodium alum Na₂SO₄ * A1₂(SO₄)₃. 24H₂O

(iii) Ammonium alum (NH₄)₂SO₄ * AI₂(SO₄)₃ 24H₂O

(iv) Ferric alum (NH₄)₂SO₄ * Fe₂(SO₄)₃ 24H₂O

Potash alum is prepared in the laboratory by mixing hot equimolar quantities of K₂SO₄ and Al₂(SO₄)₃. The resulting solution on concentration and crystallisation gives potash alum.

Note 1. A mixture of Al powder NH₄NO₃ is called ammonal and is lUed in bombs.

2. Al is the chief constituent of silver paints.

3. A1₂(SO₄)₃ 1.8 used for making fire proof clothes.

Group 14

General Physical Properties of Group 14 Elements

(i) Electronic configuration Their valence shell electronic configuration is ns² np²

(ii) Metallic character C and Si are non-metals, Ge is a metalloid and Sn and Pb are metals.

(iii) Appearance C is black. Si is light-brown, Ge is greyish, Sn and Pb are silvery white.

(iv) Density Density increases with increase m atomic number due to increase in mass per unit volume down the group.

(v) Melting points and boiling points The melting points and boiling points decrease from carbon to lead but carbon and silicon have very high melting and boiling points due to their giant structure.

(vi) Oxidation state They exhibit +2 and +4 oxidation state. The compounds of Pb in +4 oxidation state are powerful oxidising agents since, +2 oxidation state of Pb is more stable due to inert pair effect.

The compounds in +2 oxidation state are ionic in nature and in + 4 oxidation state are covalent in nature (According to Fajan’s rule).

(vii) Ionisation enthalpy It decreases from C to So. For Pb. it ie slightly higher than Sn.

(viii) Electronegativity values The value decreases from C to Pb but not in a regular manner probably due to filling of d-orbitals III and Sn and f- orbitals In Pb.

(ix) Catenation The greater the strength of element-element bond. the greater is the strength of catenation.

C >> Si > Ge = Sn > Pb (catenation).

(x) Allotropy All the elements of this group except Pb exhibit allotropy.

In cold countries white tin changes to grey tin and results in decrease in density. This is called tin disease or tin plague.

(xi) Valency All elements exhibit tetra valency. In case of carbon, 406 kJ mol^-1 of energy is required for promotion of 2s – electron to 2p.

Formation of two extra bonds provide this energy.

(xii) Atomic and ionic radii Both increase from C to Pb.

(xiii) Multiple bonding Carbon forms pπ – pπ bonds with itself and with S, N and O. Other clements show negligible tendency of this
type due to their large size. Others form dπ – pπ multiple bonds.

Chemical Properties of Group 14 Elements

(a) Hydrides All members of the group form covalent hydrides. Their number and ease of formation decreases down the group.

Hydrides of carbon are called hydrocarbons (alkanes, alkenes or alkynes).

Hydrides of Si and Ge are known as silanes and germanes.

The only hydrides of Sn and Pb are SnH4 (stannane) and PbH₄ (plumbane),

Their thermal stability decrease down the group.

Their reducing character increases down the group.

(ii) Halides All the elements give tetrahedral and covalent halides of the type MX₄ except PbBr₄, and PbI₄.

Thermal stability

CX₄ > SiX₄ > GeX₄ > SnX₄ > PbX₄

Order of thermal stability with common metals

MF₄ > MCl₄ > MBr₄ > MI₄

Except CX₄ other tetrahalides can hydrolysed due to the presence of vacant d-orbitals.

SiX₄ + 2H₂O → SiO₂ + 4Hx.

ease of hydrolysis: SiX₄ > GeX₄ > SnX₄ > PbX₄

Except C, other elements form dihalides of the type MX₂ which are nlOre ionic and have higher melting points and boiling points, e.g., SnC1₂ is a solid whereas SnC1₄ is a liquid at room temperature.

SnCl₂ . 5H₂O is called bitter of tin and is used as a mordant in dyeing.

(iii) Oxides They form two types of oxides. mono-oxides of the type MO. e.g.,

CO (neutral) and SiO, GeO. SnO. PbO(all basic) and dioxides of the type MO₂

CO₂ is linear gas at ordinary temperature. Solid CO₂ is known as dry ice or drikold.

SiO₂ is a solid with three dimensional network in which Si is bonded to four oxygen atoms tetrahedrally and covalently. A mass of hydrated silica (SiO₂) formed from skeletons of minute plants, known as diatoms, is called kieselguhr. It is a highly parous material and is used in the manufacture of dynamite.

Carbon

Free states (diamond. graphite, coal etc.) and combined states (oxides, carbonates, hydrocarbons etc.)

Allotropic Forms of Carbon

The crystalline forms include

(i) Diamond It is the hardest and has three dimensional polymeric structure in which hybridisation of C is sp³. It is covalent solid. melting point 3650°C. density 3.51 g/cm³ and bad conductor of heat and electricity.

(ii) Graphite It is dark grey. having hexagonal plates, hybridisation of each C is sp². It is good conductor of heat and electricity due to the presence of free electrons. It was also known as black lead. It is a very good lubricant.

Aqua dag Suspensions of graphite in water.

Oil dag Suspension of graphite in oil lubricants.

(iii) Fullerenes These are the only pure form of carbon. C₆₀ molecule contains 12 five membered rings and 20 six membered rings. The five membered rings are connected to six membered rings while six membered rings are connected to both five and six membered rings. These are used in microscopic ball bearings, light weight batteries, in synthesis of new plastics and new drugs.

Amorphous forms of carbon are

(i) Coal The different forms of coal are peat (60 % C), lignite (70 % C), Bituminous (78 % C), Semi Bituminous (83 % C) and anthracite (90 % C). Bituminous is most common variety of coal.

(ii) Coke It is obtained by destructive distillation of coal

(iii) Charcoal or wood charcoal It is obtained by heating wood strongly in absence of air. When heated with steam, it becomes more activated. It is used to remove colouring matters and odoriferous
gases.

(iv) Bone black or animal charcoal It is obtained by destructive distillation of bones in iron retort. By products are bone oil or pyridine. It is used as adsorbent. On burning, it gives bone ash which is calcium phosphate and used in the manufacture of phosphorous and phosphoric acid.

(v) Lamp-black It is obtained by burning vegetable oils in limited supply of air. It is used in the manufacture of printing ink, black paint, varnish and carbon paper.

(vi) Carbon-black It is obtained by burning natural gas in limited supply of air. It is added to rubber mixture for making automobile tyres.

Coal Gas

Preparation By destructive distillation of coal.

Composition

H₂ = 45 – 55 % N₂ = 2 – 12 %
CH₄ = 25 – 35 % CO₂ = 0 – 3 %
CO = 4 – 11 % O₂ = 1 – 1.5 %

Ethylene, acetylene, benzene, etc. = 3 – 5 %

Uses It is used as illuminant, as fuel and to provide inert atmosphere in the metallurgical processes.

Natural Gas

It is found along with petroleum below the surface of earth.

Composition CH₄ = 60 – 80 %

Higher hydrocarbons = 2 – 12%

C₂H₆ = 5 – 10 %, C₃H₈ = 3 – 18 %

Uses It is used as a fuel.Its partial combustion yields carbon black (reinforcing agent for rubber).

Oil Gas

Preparation

Uses It is used as fuel in laboratories in Bunsen burners.

Wood Gas

Preparation Destructive distillation of wood gives wood gas (CH₄, C₂H₆ H₂)

Uses It is used as fuel.

Liquified Petroleum Gas (LPG)

Composition n-butane + Iso-butane

Uses It is used as domestic fuel.

Carbon Monoxide (CO)

Preparation

Properties

It is colourless, odourless and almost water insoluble gas. It is a powerful reducing agent. CO is used in the extraction of many metals from their oxide ores.

Carbon Dioxide (CO₂)

Preparation

Compounds of Silicon

Silicates

Silicates are metal derivatives of silicic acid, H₂SiO₃ and can be obtained by fusing metal oxides or metal carbonates with sand. The basic structural unit of silicates is SiO^{– 4}₄.

Talc consists of planar sheets which can slip over one another due to weak forces of attraction, and is a constituent of talcum powder. That’s
Why talcum powder has a slippery touch.

Mica (abrak) is naturally occurring aluminium silicate [KH₂AI₃(SiO₄]₃ or KAI₃Si₃O₁₀(OH)₂.

Silicones

The linear, cyclic or cross linked polymeric compounds containing (R₂SiO) as a repeating units, are known as silicones. They are manufactured from alkyl substituted chlorosilanes.

Silicones are chemically inert, water repellent, heat resistant, good electrical insulators. These are used as lubricants (vaseline), insulators etc.

Carborundum

It is second hardest material known and has formula SiC (silicon carbide), It is used as high temperature semiconductor, in transistor diode rectifiers.

Glass

it is a transparent or translucent amorphous substance obtained by fusion of sodium carbonate (or sodium SUlphate), calcium carbonate and sand (silica). It is not a true solid, so its melting point is not sharp.

General formula of glass is Na₂O * CaO * 6SiO₂.

Coloured glasses are obtained by adding certain substance to the molten mass.

Different Varieties of Glass

Glass is attacked by HF. This property is used in the etching of glass

Compounds of Lead

Chrome yellow (PbCrO ₄ )

It is prepared by adding potassium chromate to lead chromate and is used as a yellow pigment under the name chrome yellow. On treating with alkali. it gives basic lead chromate or chrome red, PbCrO₄ * PbO.

Basic lead carbonate, Pb(OH)₂ . 2PbCO₃

It is also known as white lead and is prepared by adding sodium carbonate solution to any lead salt.

3Pb(NO₃)₂ + 3Na₂CO₃ + H₂O → Pb(OH)₂ * 2PbCO₃ + 6NaNO₃ + CO₂

It is used as white paint. The disadvantage of using white lead in paints is that it turns black by the action of H₂S of the atmosphere.

lead poisoning is called plumbosolvency which increases in the excess of nitrates, organic acids and ammonium salts.

Group 15

The 15 group of the Periodic Table consists of nitrogen. phosphorus. arsenic, antimony and bismuth. These elements are known as pnicogens and their compounds as pniconides.

Physical Properties of Group 15 Elements

(i) Electronic configuration Their valence shell electronic configuration is ns² np³

(ii) Metallic character N and P are non-metals, As and Sb are metalloids and Bi is metal.

(iii) Physical state Nitrogen is the first element after hydrogen which is diatomic gas in native form. All other elements in the group are solids.

(iv) Atomicity N2 is diatomic while others are triatomic E₄.

(V) Melting and boiling points The melting point increases from nitrogen to arsenic. The boiling points increase regularly on moving down the group.

(Vi) Density It increases down the group.

(Vii) Atomic radii It increases with increase in atomic number as we go down the group.
(viii) Allotropy All the elements (except Bi) exhibit allotropy. Nitrogens – α nitrogen, β – nitrogen.

Phosphorus – White, red, black
Arsenic – Grey, yellow, black
Antimony – Metallic yellow (explosive)

(ix) Oxidation state

Nitrogen has a wide range of oxidation states.

The stability of +3 oxidation state increases and stability of +5 oxidation state decreases on moving down the group due to inert pair effect.

(x) Ionisation enthalpy Ionisation energy of nitrogen is very high due to its small size and half-filled highly stable configuration. The ionisation energy decreases down the group.

(xi) Electronegativity It decreases from nitrogen to bismuth.

(xii) Catenation ‘They exhibit the property of catenation but to lesser extent due to weak E – E bond than 14 group elements.

(xiii) Reactivity Elemental nitrogen is highly unreactive because of its strong triple bond. (almost as inert as noble gases).

White phosphorus is extremely reactive and kept in water. It is inflammable and can be ignited at 45°C.

Chemical Properties of Group 15 Elements

(i) Hydrides All the elements of this group form hydrides of the type EH₃, which are covalent and pyramidal in shape. Some properties follows the order as mentioned

[These properties are

1. Thermal stability,
2. Basic strength,
3. Solubility in water,
4. Bond angle NH₃ (107.4°); PH₃ (92°),AsH₃ (91° ), SbH₃(90° ),
5. Strength of M – H bond

Some properties follow the order

NH₃ < PH₃ < AsH₃ < SbH₃ < BiH₃

[These properties are

1. Reducing character
2. Covalent character
3. Rate of combustion

(ii) Halides All the elements of this group form trihalides, MX₃ and except nitrogen all form pentahalides, MX₅, e.g., NCi₃, NI₃, PCI₃, BiCI₃, AsCI₃ , PCl₅ etc. Trihalides (except of N) behaves as Lewis acid and the order of their strength is PCl₃ > AsCl₃ > SbCl₃ Trihalides of N behave as Lewis base and has the following order of strength

NF₃ < NCl₃ < NBr₃ < NI₃

NCl₃ is an explosive compound

(iii) Oxides All the elements of this group form oxides of the type M₂O₃ and M₂O₅.

The acidic strength of pentoxide and trioxides decrease on moving down the group, i.e.,

N₂O₅ > P₂O₅ > As₂O₅ > Sb₂O₅

BiOCI is called pearl white.

Nitrogen and its Compounds

1. Dinitrogen (N2)

Preparation

Properties

1. Nitrogen does not react with alkali metals except Li but reacts with alkaline earth metals to give metal nitride.

2. Reaction with oxygen

3. Reaction with non-metals

4. Reaction with CaC₂

Uses Liquid N₂ is used as refrigerant. Nz is used in the manufacture of HNO₂, NH₂, CaCN₂(calcium cyanamide) and other nitrogenous compounds. It is used for filling electric bulbs.

2. Ammonia (NH3)

Preparation

(i) Lab method

2NH₄Cl + Ca(OH)₂ → CaCI₂ + 2NH₃ + 2H₂O

(ii) Haber’s process

Properties

1. It is a colourless gas with characteristic pungent odour. It is extremely soluble in water due to H – bonding.

2. It is a strong Lewis base and used in the metal ion detection as

3. Reaction with chlorine

When NH₃ is in excess, N₂ is the main product.

8NH₃ + 3C1₂ → 6NH₄CI + N₃

When C1₂ is in excess, NCl₃ is the main product.

NH₃ + 3C1₂ → NCl₃ + 3HCl

4. Reaction with Nesseler’s reagent

Uses It is used as a refrigerant and to produce various nitrogenous fertilizers.

Oxides of Nitrogen

NO₂ contains odd number of valence electrons. On dimerisation. it is converted to stable N₂O₄ molecule with even number of electrons .

3.Nitric acid (HNO₃)

It is a stronger acid than H₃PO₄.

Preparations

(i) Lab method

NaNO₃ + H₂SO₄ (cone.) → NaHSO₄ + HNO₃

(ii) Ostwald’s process

Physical properties It is a syrupy, colourless, pungent liquid usually available as 68 % and 15.7 M aqueous solution is often yellow due to small concentrations of NO₂.

Chemical reactions

1. Action of nitric acid on zinc under different conditions

2. Action of nitric acid on copper under different conditions

3. Reaction with non-metals

4. Brown ring test of nitrate

5. Metals like Fe. Cr. Ni, AI or Co becomes inactive or passive due to stable oxide layers.

Structure of nitric acid

Uses It is used

1. in the manufacturing of fertilizers.
2. for purification of silver and gold.
3. in the manufacturing of explosives and as oxidising agent.
4. as nitrating reagent

Phosphorus and its Compounds Allotropic Forms of Phosphorus

(i) White phosphorus
(ii) Red phosphorus
(iii) Black phosphorus

Some Points of Distinction Between White and Red Phosphorus

Black phosphorus is formed when red phosphorus is heated in a sealed tube at 803 K. It does not oxidise in air.

Match box side contains red P or P₂S₃ + glue and on tip of match stick. red P, KelO₃ chalk and glue is deposited.

Chemical properties

1. With non-metals

2. With compounds

Uses It is used in match boxes, explosives, as rat poison, in fertilizers and alloys

1. Phosphine (PH₃)

Preparation It is prepared by following methods

Properties

1. It is a colourless gas with rotten fish like smell and is highly poisonous. It explodes in contact. with traces of oxidising agents like HNO₃,C1₂ and Br₂ vapours.

3CuSO₄ + 2PH₃ → CU₃P₂ + 3H₂SO₄

3HgCl₂ + 2PH₃ → Hg₃P₂ + 6HCI

2. Phosphine is weakly basic.

PH₃ + HBr → PH⁺₄Br^–

Uses It is used to prepare smoke screens in warfare. A mixture of CaC₂ and Ca₃P₂ is used in Holme’s signals.

2. Phosphorus Trichloride (PCl₅)

Preparation

Properties It is a colourless oily liquid. having pyramidal shape [sp³ – hybridised]

3.Phosphorus Pentachloride (PCI5)

Preparation

P₄ + 10 Cl₂ → 4 PCl₅

P₄ + 10 SO₂CI₂ → 4PCl₅ + 10 SO₂

Structure PCl₅ in gaseous and liquid phases has sp³d – hybridization and its shape is trigonal bipyramidal. The three equatorial P – CI bonds are equivalent while the two axial bonds are longer equatorial bonds.

Properties In solid state, PCI₅ exists as an ionic solid, [PCI₄]⁺ [PCl₆]^– in which, the cation, [PCI₄]⁺ is tetrahedral and the anion [PCl₆]^– is octahedral.

Oxoacids of Phosphorus

In toothpaste, CaHPO₄ * 2H₂O is added as mild abrasive and polish agent.

Group 16

The elements oxygen (0), sulphur (S), selenium (Se), tellurium (Te) and polonium (Po) belong to group 16 of the Periodic Table. These elements are known as chalco gens, i.e., ore forming elements.

The name sulphur has been derived from sanskrit word ‘Sulvezi’ meaning ‘killer of copper’.

General Physical Properties of Group 16 Elements

(i) Electronic configuration Their valence shell electronic configuration is ns², np⁴.

(ii) Metallic and non-metallic character

(iii) Abundance O > S > Se > Te > Po

(iv) Density It increases down the group regularly,

(v) Melting point and boiling point Both show a regular increase down the group due to increase in molecular weight and van der Waals’ forces of attraction.

(vi) Oxidation state

In OF₂, the oxidation state of oxygen is +2.

(vii) Ionisation energy They possess a large amount of ionisation energy which decreases gradually from 0 to Po due to increase in size of atoms and increase in screening effect.

(viii) Electron affinity They have high electron affinity which decrease from O to Po. As the size of the atom increases. the extra added electron feels lesser attraction by nucleus and hence, electron affinity decreases.

(ix) Electronegativity It decreases down the group due to decrease in effective nuclear charge down the group.

(x) Catenation 16 group elements follow the order as shown below

S-S > Se-Se > O-O > Te-Te

(xi) Atomicity Oxygen is diatomic, sulphur and selenium are octaatomic with puckered ring structure.

(xii) Allotropy

Oxygen – Dioxygen (O₂) and ozone (O₃)

Sulphur – Rhombic (01′ a) sulphur. Ss

Monoclinic (or β) sulphur, S₈(most stable), plastic sulphur

(xiii) Atomic radii and ionic radii They increase regularly from O to Po.

Chemical Properties of 16 Group Elements

(i) Hydrides All these elements form stable hydrides of the type H₂E. (Where. E = 0, S, Se, Te and Po).

2H₂ + O₂ ⇔ 2H₂O

FeS + H₂SO₄ → H₂S + FeSO₄

H₂O is a liquid due to hydrogen bonding. While others are colourless gases with unpleasant smell.

[Down the group acidic character increases from H₂O to H₂Se. All the hydrides except water possess reducing property and this character increases from H₂ S to H₂ Te.

(ii) Halides The stability of the halides decreases in the order

F^– > Cl^– > Br^– > 1^–

Amongst hexahalides, hexafluorides are the only stable halides. AD hexafluorides are gaseous in nature. SF₆ is exceptionally stable for steric reasons.

SF₄ is a gas, SeF₄ is a liquid and TeF₄ is a solid. These fluoride have sp₃ d-hybridisation and see-saw geometry. They behave Lewis acid as well as Lewis base e.g.,

SF₄ + BF₃ → SF₄ → BF₃

SeF₄ + 2F^– → [SeF₆]^2-

The well known mono halides are dimeric in nature. Example are S₂F₂, S₂C1₂, S₂Br₂, Se₂C1₂ and Se₂Br₂. These dimeric halides undergo disproportionation as given below

2 SeCI₂ → SeCI₄ + Se

(ill) Oxides They form AO₂ and AO₃ type oxides. Their acidic nature follow the order

SO₂ > SeO₂ > TeO₂ > PoO₂ and SO₃ > SeO₃ > TeO₃

Ozone is considered as oxides of oxygen.

SO₂ is a gas having sps -hybridisation and V-shape.
SO₃ is a gas which is sp2-hybridised and planar in nature.
SeO₂ is a volatile solid consists of non-planar infinite chains.
SeO₃ has tetrameric cyclic structure in solid state. SO₂ and SO₃ are the anhydrides of sulphurous (H₂SO₃) and sulphuric acid (H₂SO₄) respectively.

Note In photocopying (xerox) machines Se acts as photoconductor.

Oxygen and its Compounds

1. Dioxygen

Priestley and Scheele prepared oxygen by heating suitable oxygen compounds.

Preparation By action of heat on oxygen rich compounds

(i) From oxides

(ii) From peroxides and other oxides

(iii) From certain compounds

Physical properties It is colourless, odourless, tasteless, slightly heavier than air and sparingly soluble in water.

Chemical properties On heating it combines directly with metals and non-metals, e.g.,

2Mg + O₂ → 2MgO

4Na + O₂ → 2 Na₂O → Na₂O₂

Combination with O₂ is accelerated by using catalyst. Platinum is particularly an active catalyst.

Uses It is used in welding and cutting oxy-hydrogen or oxy-acetylene torch and in iron and steel industry to increase the content of blast in the Bessemer and open hearth process. It is also used for life support systems e.g., in hospitals, for divers, miners and mountaineers.

Tests

1. With NO it gives reddish brown fumes of NO₂.

2. It is adsorbed by alkaline pyrogallol.

2. Ozone (O₃)

Preparation By passing silent electric discharge through cold, dry oxygen in ozoniser. (Lab method) –

3O₂ ⇔ 2O₃; + 284.3 kJ

Physical properties It is pale blue gas with characteristic strong smell. It is slightly soluble in water.

Chemical reactions

1.Decomposition

2. Oxidising action

3. It acts as a powerful oxidising agent. It liberates iodine from neutral KI solution and the liberated L,turns starch paper blue.

2KI + H₂ + O₃ → 2KOH + I₂ + O₂

I₂ + Starch → Blue colour

Uses It is used

1. as a germicide and disinfectant for sterilizing water.
2. ail a bleaching agent for oils, ivory wax and delicate fibres.
3. for detecting ‘the position of double bond in unsaturated compounds.
4. in destroying odours coming from cold storage room, slaughter houses and kitchen of hotels.

compounds of Sulphur

1.Sulphur Dioxide (SO₂)

Method of preparation

(i) By heating sulphur in air

Chemical reactions It turns lime water milky due to the formation of calcium bisulphite. However, in excess of SO₂ milkiness disappears due to the formation of calcium bisulphite.

when H₂S gas is passed through a saturated solution of SO₂ till its smell disappears, it turns in a milky solution the Wacken roder’s liquid. When H₂S is passed through H₂SO₄ the reaction is called Wacken roder’s reaction.

Oxoacids of Sulphur

2. Sulphuric Acid (H₂SO₄)

Sulphuric acid is one of the most important industrial chemicals world wide. It is called the king of chemicals. It is manufactured by lead chamber process or contact process. Contact process involves three steps:

(i) Burning of sulphur or sulphur ores ill air to generate SO₂.

(ii) Conversion of ₂ to ₂ by the reaction with oxygen in the presence of a catalyst (V₂O₅).

(iii) Absorption of SO₃ in H₂SO₄ to give oleum (H₂S₂O₇) which upon
hydrolysis gives H₂SO₄.

Properties

1. Sulphuric acid is a colourless, dense, oily liquid.

MX + H₂SO₄ → 2HX + M₂SO₄

2. Concentrated sulphuric acid is a strong dehydrating agent.

The burning sensation of concentrated H₂SO₄ on skin.

3. Hot concentrated sulphuric acid is a moderately strong oxidising agent. In this respect, it is intermediate between phosphoric acid and nitric acid.

Uses It is used in petroleum refining, in pigments paints and in detergents manufacturing.

3. Hypo

It is chemically sodium thiosulphate pentahydrate, Na₂S₂O₃ * 5H₂O.

Preparation 1. It is prepared by boiling sodium sulphite solution with flowers of sulphur and stirring till the alkaline reaction has disappeared.

Na₂SO₃ + S → Na₂S₂O₃

2. It is also prepared by spring’s reaction.

Na₂S + Na₂SO₃ + I₂ → Na₂S₂O₃ +2NaI

Properties 1. It is a colourless, crystalline and efflorescent substance.

2. It gives white ppt with a dilute solution of AgNO₃. Which quickly changes into black due to the formation of Ag₂S.

Uses

1. Due to its property of dissolving silver halide, it is used in photography for fixing under the name hypo.

2 Na₂S₂O₃ + AgBr → Na₃ [ Ag(S₂O₃)₂] + NaBr

2. During bleaching, it is used as antichlor,

Na₂S₂O₃ + CI₂ + H₂O → Na₂SO₄ + S + 2HCI

3. It is used to remove iodine stain, for volumetric estimation of iodine and in medicines.

Group 17

The 17 group of Periodic Table contains five elements fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (As) combinedly known as halogens (salt forming elements). Astatine is artificially prepared radioactive element.

General Physical Properties of Group 17 Elements

(i) Electronic configuration Their valence shell electronic configuration is ns², np⁵

(ii) Physical state Intermolecular forces in halogens are weak and increase down the group. Thus, F₂ and Cl₂ are gases, Br₂ is volatile liquid and I₂ is solid.

(iii) Atomicity All are diatomic in nature.

(iv) Abundance Being very reactive in nature, they are not found free in nature. Their presence in earth’s crust follows the order.

F₂ > Cl₂ > Br₂ > I₂

(v) Colour They absorb light in the visible range forming excited states and are thus, coloured in nature.

(vi) Metallic character All the elements are non-metals and metallic character increases down the group. Thus, 1 forms 1+.

(vii) Oxidation state

(viii) Bond energy and bond length The bond length increases from fluorine to iodine and in the same order bond energy decreases However, the bond dissociation energy of F2 is lesser due to its smaller size. The order of bond energy is

(he) Density It increases down the group in a regular fashion and follows the order F > Cl > Br > 1.

(x) Ionisation enthalpy The ionisation enthalpy of halogens is very high and decreases down the group. The iodine also forms I⁺ and I³⁺ and forms compounds like leI, lCN, IPO₄. In molten state, the compounds conduct electricity showing ionic character.

(xi) Electron affinity The halogens have the high values for electron affinity. The order of electron affinity is

C1₂ > F₂ > Br₂ > I₂

Due to small size of fluorine (hence, high electron density), the extra electron to be added feels more electron-electron repulsion. Therefore. fluorine has less value for electron affinity than chlorine.

(xii) Reduction potentials and oxidising nature E°_red of halogens are positive and decrease from F to I. Therefore, halogens act as strong oxidising agents and their oxidising power decreases from fluorine to iodine. Fluorine is the strongest oxidising agent and is most reactive. That’s why it is prepared by the electrolysis of a mixture of KHF₂ and anhydrous HF using Monel metal as a catalyst.

(xii) Solubility Halogens are soluble in water which follows the order

F₂ > C1₂ > Br₂ > I₂

The solubility of iodine in water is enhanced in the presence of KI.

KI + I₂ ⇔ KI₃ ⇔ K⁺ + I^–₂

I₂ forms blue colour complex with starch.

Chemical Properties of Group 17 Elements

(i) Hydrides HF is a low boiling liquid due to intermolecular hydrogen bonding, while HCI, HBr, HI are gases. The boiling point follows the trend

HF > Hi > HBr > HCl

Some other properties show the following trend :

(ii) Oxides Fluorine forms two oxides, OF₂ and O₂F₂, but only OF₂ is thermally stable at 2.98K O₂F₂ oxidises plutonium to PuF₆ and the reaction is used for removing plutonium as PuF₆ from spent nuclear fuel.

Chlorine forms a number of oxides such as, CI₂O, CI₂O₃, Cl₂O₅ , Cl₂O₇ , CIO₂ and CIO₂ is used as a bleaching agent for paper pulp, textiles and in water treatment.

Br₂O BrO₂ BrO₃ are the least stable bromine oxides and exist only at low temperatures. They are very powerful oxidising agents.

The iodine oxides, i:e., I₂O₄, I₂O₅,I₂O₇ are insoluble solids and decompose on heating. I₂O₅ is a very good oxidising agent and is used in the estimation of carbon monoxide.

(iii) Reaction with alkali

Other halogens form hypohalite with dilute NaOH and halate with cone. NaOH₄.

(iv) Oxoacids of halogens Higher oxoacids of fluorine such as HFO₂, HFO₃ do not exist because fluorine. is most electronegative
and has absence of d-orbitals.

+3 oxidation state of bromine and iodine are unstable due to inert pair effect. therefore, HBrO₂ and HIO₂. do not exist.

Acidic character of oxoacids decreases as the electronegativity of halogen atom decreases. Thus, the order of acidic strength.

For the oxoacids of same halogens. acidic strength and thermal stability increase as the number of O atoms increases

Interhalogen Compounds

When two different halogens react with each other, interhalogen compounds are formed. These compounds are covalent and diamagnetic in nature. They are volatile solids or liquids except elf which is a gas at 298 K. Interhalogen compounds are more reactive than halogens (except fluorine).

The XY₃ type compounds have bent ‘T’ shape, XY₅ type compounds have square pyramidal shape and IF₇ has pentagonal bipyramidal structure.

BrF₃ has “T” shaped structure due to 3 bp and 2 lp.

ICI is more reactive than I₂ due to weak bond. ClF₃ and BrF₃ are used for the production of UF₆ in the enrichment of ²³⁵ U.

U(s) + 3CIF₃(I) → UF₆(g) + 3CIF(g)

Pseudohalogens and Pseudohalides

The substances behaving like halogens are known as pseudohalides. Some examples are

Chlorine and its Compounds

Occurrence

Common salt, NaCI is most important. Chlorine is also present in sea water and as rock salt.

Preparation of Chlorine

Properties

It is yellowish green gas, collected by upward displacement of air poisonous in nature, soluble in water. It’s aqueous solution is known as chlorine water.

Chemical Reactions

(i) Action of water

Coloured matter + [0] → colourless matter.

The bleaching action of chlorine is due to oxidation and is permanent.

(x) Chromyl chloride test When a mixture of chloride and solid K₂Cr₂O₇ is heated with concentrated H₂SO₄ in a dry test tube, deep red vapours of chromyl chloride are evolved.

When these vapours are passed through NaOH solution, the solution becomes yellow due to the formation of sodium chromate

The yellow solution is neutralised with acetic acid and on addition of lead acetate gives a yellow precipitate of lead chromate.

Uses

It is used as a bleaching agent, disinfectant and in the manufacture of CHCl₃,CCl₄, DDT, anti-knocking compounds and bleaching powder.

Hydrochloric Acid (Hel)

Preparation

Properties

It is a colourless and pungent smelling gas. It is extremely soluble in water and ionises as below

Noble metals like gold, platinum can dissolve in aqua-regia [three part cone HCl and one part of cone HNO₃].

Uses

It is used in the manufacture of chlorides. chlorine, in textile and dyeing industries, in medicine and in extraction of glue from animal tissues and bones.

Iodine (I₂)

It’s major SOurce is deep sea weeds of laminaria variety. Their ashes which is called kelp contain 0.5% iodine as iodides.

Another source of 12 is caliche or crude chile saltpetre (NaNO₃) which contains 0.2%, NaIO₃

Iodine is purified by sublimation.

It shows no reaction with water. Tincture of iodine is a mixture of I₂ and Kl dissolved in rectified spirit.

18 Group

The 18 group of the Periodic Table consists of colourless, odourless gases at room temperature, isolated by William Ramsay in 1898 from air

General/Physical Characteristics of Group 18 Elements

(i) Electronic configuration Their valence shell electronic configuration is ns², np² except He.

(ii) Physical state They are all gases under ordinary conditions of temperature and pressure.

(iii) Abundance In 1.O% air, the abundance follows the order

Ax > Ne > He > Kr > Xe

(iv) Atomicity The Cp / Cv = 1.67 shows their monoatomic nature.

However under high energy conditions, several molecular ions such as He⁺₂, HeH⁺, HeH²⁺and Ar^{+₂ are formed in discharge tubes. They only survive momentarily and are detected spectroscopically.}

(v) Melting and boiling points Due to the increase in magnitude of van der Waals’ forces, the melting point and boiling point increases from He to Rn.

(vi) Atomic radii The atomic radii increases from He to Rn. It corresponds to the van der Waals’ radii. So it has greatest atomic size in respective period.

(vii) Density The density of noble gases increases down the group.

(viii) Heat of vaporisation They have very low values of heal of vaporisation due to weak van der Waals’ forces of attraction. The value increases down the group.

(ix) Solubility in water They are slightly soluble in water and solubility increases from He to Rn.

(x) Liqnefication It is extremely difficult to liquify inert gases due to weak van der Waals’ forces of attraction among their molecules. Hence, they posses low value of critical temperature also.

(xi) Ionisation energy All noble gases possess very stable (ns² and ns² np6) electronic configuration. Therefore. ionisation energy of noble gases is very high and decreases down the group.

(xii) Electron affinity Due to the presence of stable electronic configuration, they have no tendency to accept additional electron. Therefore, electron affinity is almost zero.

Chemical Properties of Group 18 Elements

The noble gases are inert in nature because of their completely filled subshells. In 1962, the first compound of noble gases was prepared. It is hexafluoroplatinate (prepared by Bartlett).

Xe + PtF₆ → Xe[PtF₆]

Now, many compounds of Xe and Kr are known with fluorine and oxygen.

Preparation of Compounds of Xenon

Chemical Reactions of Xenon Compounds

As the s-orbital can accommodate only two electrons, two groups (1 and 2) belong to the s-block,

The general electronic configuration of s-block elements is ns^{l or 2}

Alkali Metals [Group-I]

Group-I elements have one electron in their valence shell. They do not occur in the native or free state. These elements are collectively known as alkali metals because their oxides and hydroxides form strong alkalies like NaOH, KOH, etc. Lithium is known as bridge element.

General Characteristics of Alkali Metals

(i) Electronic configuration [noble gas] ns¹

(ii) Atomic radii The alkali metals have the biggest atomic radii in their respective periods.

Atomic radii increases as we go down the group due to the addition of a new shell in each subsequent step.

All of these have bee lattice with coordination number 4.

(iii) Ionic radii Ionic radii of the alkali metals are much smaller than their corresponding metals due to lesser number of shells and contractive effect of the increased nuclear charge.

The ionic radii of all these alkali metal ions go on increasing on moving down the group.

(iv) Density These are light metals with low densities. Lithium is the lightest known metal. On moving down the group, ‘density increases from Li to Cs.

This is because, down the group, both the atomic size and atomic mass increases but the effect of increase in atomic mass is more as compared to increase in atomic size.

The density of potassium is lesser than that of sodium because of the I abnormal increase in size on moving down from Na to K.

(v) Melting and boiling points

1. The melting and boiling points of alkali metals are quite low and decrease down the group due to weakening of metallic bond.
2. Fr is a liquid at room temperature.

(vi) Softness These are soft. malleable and ductile solids which can be cut with knife. They possess metallic lustre when freshly cut due to oscillation of electrons.

(vii) Atomic volume Atomic volume of alkali metals is the highest in each period and goes on increasing down the group from top to bottom [Li to Cs].

(viii) Ionisation enthalpy The first ionisation enthalpy of alkali metals is the lowest amongst the elements in their respective periods and decreases on moving down the group.

The second ionisation enthalpies of all the alkali metals are very high because by releasing an electron, ions acquire noble gas configuration. so removal of second electron is difficult.

(ix) Electropositive character Due to low ionisation enthalpies. alkali metals are strongly electropositive or metallic in nature and electropositive nature increases from Li to Cs due to decrease in ionization enthalpy.

(x) Oxidation state The alkali metal atoms show only +1 oxidation state, because their unipositive ions attain the stable noble gas configuration.

The alkali metal ions attain noble gas configuration with no unpaired electrons so. they are diamagnetic in nature. Alkali metals however have paramagnetic nature due to one unpaired electron.

(xi) Hydration of ions The degree of hydration depends upon the size of the cation. Smaller the size of a cation. greater is its
hydration enthalpy

Relative degree of hydration,

Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺

(xii) Flame colouration Alkali metals and their salts impart characteristic colours to the flame because the outer electrons get excited to higher energy levels, When the electron return to the original state. it releases visible light of characteristic wavelength which provides a colour to the flame.

(xiii) Photoelectric effect Due to very low ionisation enthalpy, alkali metals specially ‘Cs’ exhibit photoelectric effect ti.e., eject electrons
when exposed to light) so it is used in photoelectric cells.

(xiv) Electrical conductivity Due to the presence of loosely held valence electrons which are free to move throughout the metal structure. the alkali metals are good conductors of heat and electricity. Electrical conductivity increases from top to bottom in the order

Li⁺ < Na⁺ < K⁺ < Rb⁺ < Cs⁺

(xv) Reducing character All the alkali metals are good reducing agents due to their low ionisation energies. Their reducing character. follows the order

Na < K < Rb < Cs < Li

Chemical Properties of Alkali Metals

(i) Action of air On exposure to moist air, their surface get tarnished due to the formation of their oxides. hydroxides and carbonates.

Hence. they are kept under inert liquid like kerosene oil but lithium is kept wrapped in paraffin wax because it floats on the surface of kerosene oil due to its low density.

Note Fire due to alkali Metals is extinguished by CCI₄

(ii) Action of oxygen

(a) All the alkali metals when heated with oxygen form different types of oxides. e.g., lithium forms lithium oxide (Li₂O), sodium forms sodium peroxide (Na₂O₂), while K, Rb and Cs form superoxides MO₂ (where, M = K, Rb or Cs)

The stability of peroxides and superoxides increases as the size of alkali metal increases.

(b) Superoxides are coloured and paramagnetic as these possess three electron bondwhere one unpaired electron is present.

(c) All oxides. peroxides and superoxides are basic in nature.

Basic strength of oxides increase in the order

Li₂O < Na₂O < K₂O < Cs₂O

Na₂O₂ acquires yellow colour due to the presence of superoxides as an impurity.

KO₂ (potassium superoxide) is used as a source of oxygen in submarines, space shuttles and in emergency breathing apparatus such as oxygen masks.

(iii) Action of water or compounds containing acidic hydrogen

2M + 2H₂O → 2MOH + H₂ (where, M = Li, Na, K, Rb, and Cs)

The reactivity order with water is

Li < Na < K < Rb < Cs

This is due to increase in electropositive character in the same order.

KOH is stronger base than NaOH .

LiOH is used to remove carbon dioxide from exhaled air in confined quarters like submarines and space vehicles.

(iv) Action of hydrogen

2M + H₂O → 2MH (where, M = Li, Na, K, Rb, and Cs)

The reactivity of alkali metals towards hydrogen is

Li > Na > K > Rb > Cs.

(v) Reaction with halogens Alkali metals combine readily with halogens to form ionic halides M⁺X- (with the exception of some lithium halides).

2M + X₂ → 2M⁺ X^– (where, M = Li, Na, K etc., and X = F, Cl, Br, 1]

The reactivity of alkali metals towards a particular halogen increase in the order

Li < Na < K < Rb < Cs For a given halide, ionic character increases as the size of metal ion increases. LiX > NaX < KX < RbX < CSX All alkali halides except LiF are freely soluble in water (LiF is soluble in non-polar solvents because it has strong covalent bond). LiCl is more covalent than KCI due to smaller size of Li. Bigger the anion, larger is its polarisability. Hence, the covalent character follow the order LiI > LiBr > LiCl > LiF

(vi) Solubility in liquid ammonia All alkali metals dissolve in liquid ammonia giving deep blue solution due to formation of ammoniated metal cations and ammoniated electrons in the solution.

The blue colour is due to the excitation of ammoniated electron to higher energy levels and the absorption of photons occurs in the red region of the spectrum.

This solution is highly conducting and paramagnetic because of the presence of ammoniated electrons and ammoniated cations.

(vii) Nature of carbonates and bicarbonates Li₂CO₃ is unstable toward heat.

The thermal stability of carbonates increases on moving down the group as

All the bicarbonates (except LiHCO₃ which exists in solution) exist as solids and on heating form carbonates.

The solubility of the carbonates and bicarbonates increases on moving down the group due to decrease in lattice enthalpies. Thus, the order is

LiHCO₃ < NaHCO₃ <KHCO₃ < RbHCO₃ < CsHCO₃

A mixture of Na₂CO₃ and K₂CO₃ is known as fusion mixture K₂CO₃is known as pearl ash.

(viii) Nature of nitrates LiNO₃ on heating decomposes to give NO₂ and O₂, while the nitrates of the other alkali metals decompose on heating and give nitrites and O₂.

NaNO₃ is called chile saltpeter and KNO₃ is called Indian saltpeter.

(ix) Nature of sulphates Li₂SO₄ is insoluble in water whereas the other sulphates, i:e., Na₂SO₄, K₂SO₄ are soluble in water.

Na₂SO₄ . 10H₂O is called Glauber’s salt

Anomalous Behaviour of Lithium

Lithium shows anomalous behaviour due to the following reasons:

1. It has the smallest size in its group.

2. It has very high ionization enthalpy and highest electronegativity in the group.

3. Absence of d-orbitals inits valence shell.

As a result, it differs from the other alkali metals in the following properties :

• Lithium is harder than other alkali metals, due to strong metallic bond.
• Lithium combines with O2 to form lithium monoxide, Li₂O whereas other alkali metals form peroxides (M₂O₂) and superoxides (MO₂).
• Lithium, unlike the other alkali metals, reacts with nitrogen to form the nitride.

• Li₂CO₃ ,LiF and lithium phosphate are insoluble in water while the corresponding salts of other alkali metals are soluble in water.
• Li₂CO₃ decomposes on heating to evolve CO₂ whereas other alkali metal carbonates do not.
• Lithium nitrate on heating evolves O₂ and NO₂ and forms Li₂O while other alkali metal nitrates on heating form their respective nitrites.

Diagonal Relationship

Lithium shows diagonal resemblance with magnesium [the element of group 2] and this resemblance is due to similar polarising power, i.e.,
[ionic charge / (ionic radius)²] of both these elements.

Lithium resembles magnesium in the following respects :

1. The atomic radius of lithium is 1.31 Å while that of magnesium is 1.34 Å.
2. The ionic radius of Li⁺ⁱ on is 0.60 Å, which is very close to that of Mg²⁺ ion (0.65 Å).
3. Lithium (1.0) and magnesium (1.2) have almost similar electronegativities.
4. Both Li and Mg are hard metals.
5. LiF is partially soluble in water like MgF₂.
6. Both combine with O₂ to form monoxides, e.g., Li₂O and MgO.
7. Both LiOH and Mg(OH)₂ are weak bases.
8. Both LiCI and MgCl₂ are predominantly covalent.
9. Both Li and Mg combine with N₂ to form their respective nitrides, Li₃N and Mg₃N₂.
10. Both lithium and magnesium nitrates on heating evolve NO₂ and O₂ leaving behind their oxides.

Compounds of Sodium

1. Sodium Chloride, Common Salt or Table Salt [NaCI]

Sea water contains 2.7 to 2.9%by mass of the salt. Sodium chloride is obtained by evaporation of sea water but due to the presence of impurities like CaCl₂ and MgCl₂ it has deliquescens nature. It is purified by passing HCI gas through the impure saturated solution of
NaCl and due to common ion effect, pure NaCl gets precipitated. 28% NaCl solution is called brine.

2. Sodium Hydroxide or Caustic Soda [NaOH]

Methods of preparation

(i) A 10% solution of Na₂CO₃ is treated with milk of lime (Causticizing process).

(ii) Electrolytic process involves Nelson cell and Castner-Kellner cell.

A brine solution is electrolysed using a mercury cathode and a carbon anode. Sodium metal discharged at the cathode combines with Hg to form Na-amalgam. Chlorine gas is evolved at the anode.

The amalgam is treated with water to give sodium hydroxide and hydrogen gas.

2Na-Hg + 2H₂O → 2NaOH + 2Hg + H₂

Physical properties

Sodium hydroxide is a white translucent solid. It is readily soluble in water. Crystals of NaOH are deliquescent.

Chemical properties

1. It is a hygroscopic, deliquescent white solid, absorbs CO₂ and moisture from the atmosphere

3. Sodium Carbonate or Washing Soda (Na₂CO₃ . 10H₂O)

Solvay process

CO₂ gas is passed through a brine solution saturated with NH₃

Sodium bicarbonate is filtered and dried. It is ignited to give sodium carbonate.

Properties

1. Sodium carbonate crystallises from water as decahydrate which effloresces on exposure to dry air forming monohydrate which on heating change to anhydrous salt (soda-ash).

Uses

1. It is used in water softening, laundering and cleaning.
2. It is used in paper, paints and textile industries

4. Sodium Bicarbonate or Baking Soda (NaHCO₃) Preparation

It is obtained as an intermediate product in Solvay process.

Properties

Uses

1. It is used as a constituent of baking powder which is a mixture of sodium bicarbonate, starch and potassium bitartrate or cream of tartar and in medicine to remove acidity of the stomach (as antacid).
2. NaHCO₃ is a mild antiseptic for skin infections.
3 It is used in fire extinguisher.

5. Microcosmic salts [Na(NH₄)HPO₄ . 4H₂O]

Preparation

It is prepared by dissolving Na₂HPO₄ and NH₄Cl in the molecular proportions in hot water followed by crystallisation.

Properties

On heating it forms a transparent glassy bead of metaphosphate, which gives coloured beads of orthophosphates when heated with coloured salts like that of transition metal ions(Cu²⁺, Fe²⁺, Mn²⁺, Ni²⁺, CO²⁺).

This test is called microcosmic bead test.

Na(NH₄) HPO₄ → NH₃ + H₂O + NaPO₃

sodium metaphosphate

CUSO₄ → CuO + SO₃

CuO + NaPO₃ → CuNaPO₄ (blue bead)

It is especially used to detect silica which being insoluble in NaPO₃ gives a cloudy bead.

Alkaline Earth Metals [Group-ll]

Group-II elements are Be. Mg, Ca. Sr. Ba and Ra. which have two electrons in their valence shell. These are commonly called alkaline earth metals because their oxides are alkaline in nature and are found in earth’s crust.

Mg is present in chlorophyll and Ca is present in bones as calcium phosphate.

General Characteristics of Alkaline Earth Metals

(i) Electronic configuration [noble gas] ns²

(ii) Atomic radii and ionic radii The atomic radii and ionic radii of these clements are quite large but smaller than those of the corresponding alkali metals. due to increased nuclear charge of these elements. The atomic as well as ionic radii go on increasing down the group due to the gradual addition of extra energy level.

(iii) Density These are much denser than alkali meta Is because of their smaller size and greater nuclear charge and mass. The density. however. first decreases from Be to Ca and then steadily increases from Ca to Ra due to difference in type of crystal structure.

(iv) Melting and boiling points These metals have higher melting and boiling points than those of alkali metals because of greater number of bonding electrons.

The melting and boiling points decrease down the group with the exception of magnesium,

(v) Metallic properties These are silvery white metals. soft in nature but harder than alkali metals due to stronger metallic bonding.

(vi) Ionization enthalpy The first ionisation enthalpy of alkaline earth metals are higher than those of the corresponding alkali metals due to smaller size and ns² configuration.

The second ionisation enthalpy values are higher than their first ionisation enthalpy values but much lower than the second ionisation enthalpy values of alkali metals.

On moving down the group. due to increase in atomic size, the magnitude of ionisation enthalpy decreases.

(vii) Electropositive character These are strong electropositive elements due to their large size and comparatively low ionisation enthalpy.

On moving down the group, the electropositive character increases due to increase in atomic radii and decrease in ionisation enthalpy.

(viii) Oxidation state Alkaline earth metals uniformly show an oxidation state of +2.

In the solid state, the dipositive ions (M²⁺) form strong lattices due to their small size and high charge (i.e., high lattice enthalpy).

[In the aqueous solution, the M²⁺ cations are strongly hydrated due to their small size and high charge. The hydration energy released by the M²⁺ cation is very high]

(ix) Flame colouration Alkaline earth metals salts impart characteristic colours to the flame.

As we move down the group from Ca to Ba, the ionisation enthalpy decreases, hence the energy or the frequency of the emitted light increases. Thus,

Be and Mg because of their high ionisation energies, do not impart any characteristic colour to the flame.

(x) Crystal lattice Be and Mg crystallise in hcp, Ca and Sr in ccp and Ba in bee lattice.

Chemical Properties of Alkaline Earth Metals

Alkaline earth elements are quite reactive due to their low ionisation energies but less reactive than alkali metals. Reactivity of the group.2
elements increases on moving down the group because their ionisation enthalpy decreases.

(i) Reaction with water Group-2 elements are less reactive with water as compared to alkali metals.

M + 2H₂O → M(OH)₂ + H₂ (where, M = Mg, Ca, Sr or Ba)

Be does not react even with boiling water and Ba react vigorously even with cold water. Thus. increasing order of reactivity with water is

Mg < Ca < Sr < Ba

A suspension of Mg(OH)₂ in water is called milk of magnesia.

Ca(OH)₂ solution (lime water) and Ba(OH)₂ solution (baryta are used for the detection of CO₂,

(ii) Reaction with oxygen The affinity towards oxygen increases down the group. Thus. Be. Mg and Ca when heated with O₂ form monoxides while Sr. Ba and Ra form peroxides.

(iii) Reaction with acids Alkaline earth metals except Be, displace H₂ from acids.

Reactivity increases down the group from Mg to Ba. Only Mg displaces H₂ from a very dilute HNO₃.

(iv) Reaction with hydrogen Except Be, all other elements of group-2 combine with hydrogen on heating to form hydride (MH₂).

M + H₂ → MH₂

BeH₂ and MgH₂ are covalent and polymeric whereas the hydrides of Ca, Sr and Ba are ionic in nature.

(v) Reaction with halogens All the elements of group-2 combine with halogens at high temperature, forming their corresponding halides (MX₂).

Beryllium halides (BeF₂, BeCI₂, etc) are covalent, hygroscopic and fume in air due to hydrolysis, BeC1₂ exists as a dimer, The halides of other alkaline earth metals are fairly ionic and this character increases as the size of the metal increases.

The halides are soluble in water and their solubility decreases in the order

MgX₂ > CaX₂ > SrX₂ > BaX₂

(vi) Reaction with nitrogen These metals react with nitrogen to form nitrides of the types M₃N₂which are hydrolysed with water to evolve NH₃.

3M + N₂ → M₃N₂
M₃N₂ + 6H₂O → 3M(OH)₂ + 2NH₃

(vii) Reaction with carbon These metals when heated with carbon, form their respective carbides of the general formula MC₂ (except Be).

All these carbides are ionic in nature and react with H₂O to form acetylene (except Be₂C which gives methane).

CaC₂ + 2H₂ → Ca(OH)₂ + HC = CH

(viii) Reducing character All the alkaline earth metals are strong reducing agents because of their lower electrode potentials but these are weaker than the corresponding alkali metals.

As we move down the group from Be to Ra, the reducing character increases due to decrease in ionisation enthalpy.

(ix) Solubility in liquid ammonia Like Alkali Metals, these metals also dissolve in liquid ammonia by giving coloured solutions.

M +(x+ y)NH₃ [M(NH₃)_x]²⁺ + 2[e(NH₃)_y]^–

The tendency to form ammoniates decreases with increase in size of the metal atom (i.e., on moving down the group).

(x) Complex formation It is favoured in case of alkaline earth metals because of their small sizes as compared to the alkali metals. Both Mg²⁺ and Ca²⁺ form six coordinate complexes with EDTA (ethylenediamminetetracetic acid) which are used to determine the hardness of water.

(xi) Basic strength of oxides and hydroxides BeO and Be(OH)₂ are amphoteric while the oxides and hydroxides of other alkaline earth metals are basic. The basic strength, however, increases from Be to Ba.

The basic character of hydroxides of group – 2 elements is lesser than those of group-l hydroxides because of the larger size of later than former group.

(xii) Thermal stability and nature of bicarbonates and carbonates Bicarbonates of these metals do not exist in solid state but are known in solution only. When these solutions are heated, these get decomposed to evolve CO₂.

The carbonates of alkaline earth metals can be regarded as salts of weak carbonic acid (H₂CO₃) and metal hydroxide, M(OH)₂. The carbonates decompose on heating forming metal oxide and CO₂.

Anomalous Behaviour of Beryllium

Beryllium, differs from the rest of the members of its group due to the following reasons

1. Beryllium has a small atomic and ionic size.
2. It has no vacant d-orbitals.
3. It has a high charge density.

The points of difference are :

(i) Hardness Beryllium is denser and harder than other members of the family.

(ii) Melting point Beryllium has high melting point i.e., 1551 K while that of magnesium is 924 K

(iii) Ionisation potential It has higher ionisation potential as compared to the rest of the members of this group.

(iv) Reaction with acids Due to lower oxidation potential of Be, it does not liberate hydrogen from acids readily.

(v) Reaction with water Beryllium does not react with water even at higher temperature while other members of the family liberate hydrogen by reacting with water at room temperature.

(vi) Amphoteric in character Oxide (BeO) and hydroxide [Be(OH)₂] of beryllium are amphoteric in character and dissolve in acids to form salt and beryUate in alkali.

(vii) Formation of carbides Beryllium when heated with carbon form Be₂C which on reaction with water gives methane. While other members of the group form ionic carbide MC₂ (acetylide) which on reaction with water evolve acetylene.

Diagonal Relationship Between Be and AI

The main identical physical and chemical properties of Be with aluminium are given below

(i) Action of air Both the metals are stable in air.

(ii) Action with water Be and Al do not decompose water even at 373 K It is due to their less electropositive character.

(iii) Electropositive character Beryllium like aluminium is less electropositive due to their small ionic radii.

(iv) Complex formation Beryllium and aluminium form a number of complexes. Both form fluoro complex anions like BeF^2-₄ and AlF^3-₆ in solution.

(v) Reaction with alkali Beryllium and aluminium react with sodium hydroxide liberating hydrogen.

(vi) Passive nature Both these metals are rendered passive on reaction with concentrated nitric acid due to the formation of oxide layer on their surfaces.

(vii) Amphoteric character of oxides Oxides of both Be and Al are amphoteric in nature. So. they dissolve.both in acids as well as in alkalies.

Application of Alkaline Earth Metals and their Compounds

1. Beryllium (Be) is used in corrosion resistant alloys.
2. Alloy of Mg with aluminium is used as structural material because of its high strength. low density and ease in machining.
3. Strontium carbonate is used for the manufacture of glass for colour TV picture tubes.
4. Hydrated calcium chloride, CaCI₂ . 6H₂O is widely used for melting ice on roads, particularly in very cold countries, because a 30% eutectic mixture of CaCl₂ / H₂O freezes at -55°C as compared with NaCI / H₂O at -18°C.
5. Barium sulphate being insoluble in water and opaque to X-rays is used under the name barium meal to scan the X-ray of the human digestive system.
6. Magnesium is present in chlorophyll, a green pigment in plant, essential for photosynthesis.
7. Anhydrous CaCl₂ because of its hygroscopic nature is a good drying agent but it cannot be used to dry alcohols/ammonia/amines.
8. Magnesium perchlorate Mg(CIO₄)₂ is used as a drying agent under the name anhydrone.

Note Kidney stones generally consist of calcium oxalate. CaC₂O₄ . H₂O which dissolves in dilute strong acids but remains insoluble in bases.

Compounds of Calcium

1. Calcium Oxide or Quick Lime or Lime [CaO] Preparation

By the thermal decomposition of calcium carbonate.

Properties

1. It is a basic oxide.
2. Its aqueous suspension is known as slaked lime

3. On heating with ammonium salts, it gives ammonia.

4. It reacts with carbon to form calcium carbide

5. It is used as basic flux, for removing hardness of water, for preparing mortar (CaO + sand + water).

2. Calcium Hydroxide or Slaked Lime or Lime Water [Ca(OH)₂]

Preparation

By dissolving quicklime in water.

CaO + H₂O Ca(OH)₂;

ΔH = – 63 kj

Properties

1. Its suspension in water is known as milk of lime.

2. It gives CaCO₃ (milky) and then Ca(HCO₃)₂ with CO₂,

3. It reacts with Cl₂ to give bleaching powder, CaOCI₂

Ca(0H)₂ + Cl₂ → CaOCI₂ + H₂O

3. Calcium Carbonate or Limestone or Marble or Chalk [CaCO₃]

Preparation

By passing CO₂ through lime water.

Ca(0H)₂ + CO₂ → CaCO₃ ↓ + H₂O

Properties

It is insoluble in H₂O but dissolves in the presence of CO₂, due to the formation of calcium bicarbonate.

4. Gypsum, Calcium Sulphate Dihydrate (CaSO₄ * 2H₂O)

It is also known as alabaster.

On heating at 390 K, it gives plaster of Paris.

It is added to cement to slow down its rate of setting.

5. Plaster of Paris or Calcium Sulphate Hemihydrate (CaSO₄ * 1 / 2 H₂O)

When it is mixed with water, it forms first a plastic mass which sets into a solid mass with slight expansion due to dehydration and its reconversion into gypsum. It is obtained when gypsum is heated at 393 K.

CaSO₄ * 2H₂O → CaSO₄ * 1 / 2 H₂O + 3 / 2 H₂O

Above 393 K no water of crystallization is left and anhydrous calcium sulphate is obtained. It is known as dead burnt plaster.

6. Bleaching Powder (CaOCI₂)

It is also called calcium chloro hypochlorite or chloride of lime.

Preparation

Ca(OH)₂ + CI₂ → CaOCI₂ + H₂O

Properties

1. Its aqueous solution gives Ca²⁺, CI^– and OCl^– ions.

2. With limited quantity of dil H₂SO₄ it gives nascent oxygen which is responsible for its oxidising and bleaching action.

2CaOCI₂ + H₂SO₄ → CaC1₂ + CaSO₄ + 2HCIO

HCIO HCI + [O]

3. With excess of dil H₂SO₄ (or CO₂) it forms C1₂, which is known as available chlorine.

CaOC1₂ + H₂SO₄ → CaSO₄ + H₂O + Cl₂

CaOC1₂ + CO₂ → CaCO₃ + C1₂

The average percentage of available chlorine is 35-40%. Theoretically it should be 49~o. which diminishes on keeping the powder due to following change

6CaOC1₂ → 5CaCl₂ + Ca(CIO₃)₂

Uses It is used for bleaching, as disinfectant and germicide in sterlisation of water. for making wool unshrinkable and in the manufacture of chloroform.

7. Cement

Cement is an important building material. It is a product obtained by combining materials such as limestone (provides lime and clay provides alumina and silica, Si02 along with the oxides of iron and magnesium.) The average composition of portland cement is

CaO, 50-60%; SiO₂, 20-25%; Al₂O₃, 5-10%; MgO, 2-3%; Fe₂O₃, 1-2% and SO₃, 1-2%.

A mixture of lime (CaO) and sand in the ratio 1 : 3 with enough water to make a thick paste is called mortar.

By ash, a waste product of steel industry. has properties similar to cement and can be added to cement to reduce its cost without affecting its quality.

Position of Hydrogen in the Periodic Table

Hydrogen resembles with alkali metals (group I) as well as halogens (group 17), At the same time, it differs from both in certain characteristics. That is why hydrogen if; called “rogue element”.

However. it has been placed in group 1 on the basis of its configuration 1s¹, which is the basis of modern classification of elements.

Isotopes of Hydrogen

Hydrogen exists in the form of three sotopes :

Dihdrogen [H₂]

Methods of Preparation

(a) Lab pletbods

Metals which have reduction potential lesser than H, can liberate H₂ from acids.

Pure zinc is not used because it reacts slowly. The presence of some impurities increases the rate of reaction due to the
formation of electrochemical couples Cone sulphuric acid is also not used because it oxidises, H₂formed into H₂O.

Zn + 2H₂SO₄(conc.) → ZnSO₄ + SO₂ + 2H₂O

(ii) It can also be prepared by the reaction of zinc with aqueous alkali.

b) Commercial production of dihydrogen

(i) By the electrolysis of acidified water

ii) From water gas (Bosch process)

Carbon dioxide is removed by dissolving it in water under pressure (20-25 atm) and hydrogen left behind is collected.

(iii) From steam (Lane’s process) Super heated steam is passed over iron filings heated to about 1023-1073 K when hydrogen is
formed.

(iv) Highly pure (> 99.95%)dihydrogen is obtained by electrolysing warm aqueous barium hydroxide solution between nickel
electrodes.

(v) From hydrocarbons by partial oxidation

vi) It is also obtained as a by-product in the manufacture of NaOH and chlorine hy the electrolysis of brine solution.

During electrolysis, the reactions that take place are

Physical Properties of Dihydrogen

Dihydrogen is a colourless, odourless, tasteless, combustible gas. It is lighter than air and insoluble in water. It is neutral to litmus.

Chemical Properties of Dlhydrogen

(i) Reactivity The relative inertness of dihydrogen at room temperature is because of Its high enthalpy of H-H bond i.e.. high bond dissociation energy. So its reactions take place under specific conditions only (at high temperature).

(ii) Action with non-metals

(iii) Reaction with metals Here H₂ acts as oxidising agent.

(iv) Reducing action of dihydrogen

(v) Reactions with metal ions and metal oxides

(vi) Reaction with organic compounds

Uses of Dihydrogen

1. It is used in the manufacture of CH₃OH.
2. It produces temperature of 2850°C and oxy-atomic hydrogen flame produces a temperature of 4000°C, so it is used in oxy-hydrogen flame.
3. The largest single use of H₂ is in the synthesis of NH₃ which is used in the manufacture ofHNO₃ and fertilizers.
4. Liquid hydrogen (LH₂) is used as rocket fuel.
5. H₂ is used as a reducing agent in extraction of metals.
6. H₂ is used in fuel cell for generating electrical energy.
7. Hydrogen is used in the manufacture of synthetic petrol.

(By heating H₂ with coal and heavy oils under very high pressure in the presence of catalyst.)

Different Forms of Hydrogen

Atomic Hydrogen

It is very reactive and its half-life period is 0.33 s.

Nascent Hydrogen

Freshly prepared hydrogen is known as nascent hydrogen and is more reactive than ordinary hydrogen. It causes the reduction of certain compounds which IS not possible with ordinary hydrogen. It can never be isolated.

Activity of nascent H depends upon the reaction by which it is obtained.

Adsorbed Hydrogen

Adsorption of hydrogen at the metal surface is called occlusion. This hydrogen brings out many chemical changes such as reduction and hydrogenation. Occlusion decreases with rise in temperature.

Ortho and Para Hydrogen

When in hydrogen molecule, the nuclear spins are in the same direction, it is known as ortho hydrogen. On the other hand when the nuclear spins arc in tho opposite direction. it is known as para hydrogen. At room temperature hydrogen consists of 75% ortho and 25% para hydrogen.

Hydrogen Economy

Hydrogen economy is the use of liquid hydrogen as an alternate source of energy. The technology involves the production, transportation and storage of energy in the form of liquid hydrogen. Large scale production of hydrogen can be done by electrolysis of water or by thermochemical reaction cycle. Storage of hydrogen in liquid form can be done in vacuum insulated cryogenic tanks or in a metal or in all alloy like iron-titanium alloy as interstitial hydride. Hydrogen fuel has many advantage over conventional fuels in that it is non-polluting and it liberates large amount of energy on combustion.

Pbotohydrogen is used to obtain renewable energy from sunlight by using microscopic organism such as bacteria or algae.

Hydrides

The compounds of hydrogen with metals and non-metals are called hydrides.

Ionic Hydrides

These are formed by elements of group I, II, (except Be and Mg) by heating them in hydrogen.

These are white colourless solids (crystalline) having high m.p. and b.p. easily decomposed by water, CO₂ or SO₂.

CaH₂ + 2H₂O → Ca(OH)₂ + 2H₂

CaH₂ + 2CO₂ → (HCOO)₂Ca

They are strong reducing agents. Alkali metal hydrides are used for making LiAlH₄, NaBH₄, etc and for removing last traces of water from organic compounds.

Molecular or Covalent Hydrides

These are formed by elements of p-block having higher electronegativity than hydrogen.

1. Electron deficient hydrides These are the hydrides which do not have sufficient number of electrons needed to form normal covalent bonds, e.g., BH₃, AlH₃, etc.
2. Electron precise hydrides These are the hydrides which have exact number of electrons needed to form normal covalent bonds. e.g. hydrides of group 14 (CH₄, SiH₄, etc.)
3. Electron rich hydrides These are the hydrides which have greater number of electrons than required to form normal covalent bonds. e.g., hydrides of group 15, 16, 17, (NH₃, PH₃ ,H₂S, HF, HCl, etc). The excess electrons in these hydrides are present as lone pairs of electrons.

Metallic or Interstitial Hydrides

The transition metals and rare earth metals combine with hydrogen to from interstitial hydrides. They exhibit metallic properties and are powerful reducing agents. They are non-stoicluometric hydrides and their composition varies with temperature and pressure for e.g., LaH_2.76, TiH_1.73.

Metals of group 7, 8 and 9 do not form hydrides and this region of the Periodic Table is called hydride gap.

Polymeric Hydrides and Complex Hydrides

Polymeric hydrides are formed by’ elements having electronegativity in the range 1.4 to 2.0, e.g., (BeH₂)_n, (AlH₃)_n, etc. In complex hydrides H^– acts as ligand and is attached to central metal atom, e.g., LiAlH₄, LiBH₄, etc.

Water

Water is the most abundant and widely distributed on the earth. It occurs in all the three physical states. H₂O is a covalent molecule in which oxygen is sp³ hybridised. It has bent structure.

Physical Properties of Water

1. Water is a colourless, odourless, tasteless liquid. It has abnormally high b.p., f.p., heat of vaporisation due to hydrogea bonding.
2. Pure water is not a good conductor so it is made conductor by adding small amount of acid or alkali.
3. Density of ice (which is mass per unit volume) is lesser than that of water and it floats over water.
4. Water has maximum density at 4°0.
5. Water is a highly polar solvent with high dielectric constant 78.39. It interacts with polar or ionic substances effectively with the release of considerable amount of energy due to ion dipole interaction. The dissolution of covalent compounds like urea, glucose and C₂H₅OH, etc is due to the tendency of these molecules to form hydrogen bond with water.

Chemical Properties of Water

1. Water is amphoteric in nature.

2. In redox reactions,water reacts with metals and non- metals both.

3. In hydrated salts, water may remain in five types such as coordinated water, hydrogen bonded water, lattice water, clathrate water and zeolite water.

4. A number of compounds such as calcium hydride, calcium phosphide. etc ., undergo hydrolysis with water.

Purification of Water

It involves two processes

1. Removal of suspended impurities
2. Destroying the bacteria.

Suspended particles are removed by coagulation with alum followed by filtration.

Exposure to sunlight, boiling, chlorination (treatment with liquid Cl₂ or bleaching powder), ozonisation and addition of CuSO₅ are some processes which are employed to destroy bacteria.

Heavy Water [D₂O]

It was discovered by Urey in 1932. It can be prepared by exhaustive electrolysis of ordinary water using nickel electrodes. It is colourless, odourless, tasteless liquid.

Chemical Reactions of Heavy Water

Usesof Heavy Water

It is used

1. in nuclear reactors to slow down the speed of neutrons and called moderator.
2. as a tracer compound to study the mechanisms of many reactions.

Soft and Hard Water

The water which produces large amount of lather with soap is known as soft water and which forms a scum with soap is known as hard water.

Typesof Hardness of Water

• Temporary hardness It is due to the presence of bicarbonates of calcium and magnesium.
• Permanent hardness Tt is due to the presence of chlorides and sulphates of calcium and magnesium.

Removal of Temporary Hardness

It can be achieved:

(a) By boiling The soluble bicarbonates are converted into insoluble carbonates.

(b) By Clark’s process By adding lime water or milk of lime.

Removal of Permanent Hardness

(i) By adding washing soda The calcium or magnesium salts are precipitated as carbonates.

(ii) By adding caustic soda The temporary and permanent hardness can be removed by adding caustic soda.

(iii) By adding sodium phosphate (Na₃PO₄) The phosphates of calcium and magnesium are precipitated.

Similarly, magnesium also precipitate out in the form of magnesium phosphate, Mg₃(PO₄)₂.

(iv) Calgon’s process Calgon is sodium hexa metaphosphate (Na₆P₆O₁₈). This calgon when added to hard water form soluble complex.

Similarly. Mg²⁺ can also precipitate as Na₂[Mg₂(PO₃)₆] and water becomes free from Ca²⁺ and Mg2+ Ions.

(v) Permutit process Permutit is hydrated sodium aluminium silicate Na₂Al₂Si₂O₈.xH₂O. It exchanges its sodium ions for divalent ions Ca²⁺ and Mg^2+..

Permutit when fully exhausted can be regenerated by treating with 10% solution of sodium chloride. It is most efficient method to gel water with zero degree of hardness.

(vi) By synthetic resins

These are of two types:

(a) Cation exchange resins are big molecules containing sulphonic acid group (-SO₃H). It is first changed into sodium salt with the general formula RNa. The hard water is passed through it so Ca²⁺ and M²⁺ are exchanged and removed.

The resins like permutit can be regenerated with a solution of NaCl.

(b) Anion exchange resins are also big molecules and can exchange anions. They contain an amino group.

The water is first passed through cation resins and then through anion tesin and pure distilled water is obtained.

Measurement of Degree of Hardness

Degree of hardness is defined as the number of parts of calcium carbonate or equivalent to various calcium and magnesium salts present in one million parts of water by mass. It is expressed in ppm.

Degree of hardness (in ppm) = (wt. of CaCO₃ (g)/ wt. of hard water (g)) x 10⁶

The molecular wt. of Ca(HCO₃)₂, Mg(HCO₃)₂, CaCl₂, MgCl₂, CaSO₄ and MgSO₄ is 162, 146, 111, 95, 136 and 120 respectively. The mol. wt. of CaCO₃ is 100.

Thus, 162 g Ca(HCO₃)₂, 146 g Mg(HCO₃)₂, 111 gCaCl₂, 95 g MgCl₂ 136 g CaSO₄ and 120 g MgSO₄ are equivalent to 100 g CaCO₃.

Hydrogen Peroxide [H₂O₂]

H₂O₂ was discovered by J.L. Thenard in 1818. It is an important compound used in pollution control treatment of domestic and industrial effluents.

Methods of Preparation

Strength of Hydrogen Peroxide

The most common method to express the strength of H₂O₂ is in terms of the volume (in mL) of oxygen liberated at NTP by decomposition or 1 mL of that sample of H Z 0 2. A solution of H₂O₂labelled as ’10 volume’ actually means “1 mL of such a solution of H₂O₂ on decomposition by heat produces 10 mL of oxygen at NTP”.

(i) Strength of H₂O₂ in terms of normality

(68 x X/22.4) = 17 x N ⇒ X = 5.6 x N

where, X is volume strength of H₂O₂.

(ii) % strength = (17/56) x volume strength

(iii) X = 11.2 x molarity.

Storage of Hydrogen Peroxide (H₂O₂)

It is stored in the presence of traces of alcohol, acetanilide or sodium pyrophosphate which slow down the rate of decomposition of hydrogen peroxide.

Chemical Properties of H₂O₂

1. Acidic nature It is weakly acidic in nature and pure hydrogen peroxide turns blue litmus red.
2. Oxidising agent It acts as a strong oxidising agent in acidic as well as in basic medium.

e.g., oxidising action of HzOz is

(iii) Reducing agent

(a) In acidic medium

(b) In basic medium

(iv) Bleaching properties Its bleaching action is due to oxidation by atomic oxygen and permanent.

H₂O₂ rarr; H₂O + [O]

dye + [O] → dye is oxidised and bleached