Structures of functional groups

In alcohols, the oxygen of the -OH group is attached to sp³ hybridised carbon by a sigma bond formed by the overlap of sp³ hybrid orbital of carbon with an sp³ hybrid orbital of oxygen. The figure shown below illustrates the bonding in methanol.
The C-O-H bond angle in alcohol is slightly less than the tetrahedral angle (109⁰28). It is due to the repulsion between the unshared electron pairs of oxygen.

In phenols, the –OH group is attached to sp² hybrid carbon of an aromatic ring. The C-O-H bond angle in phenol is 109⁰. The carbon-oxygen bond length (136pm) in phenol is slightly less than that in methanol (142pm). This is due to partial double bond character on account of the conjugation of unshared electron pair of oxygen with the aromatic ring.

In ethers, the four electron pairs, ie; the two bond pairs and two lone pairs of electrons around oxygen are arranged approximately in a tetrahedral arrangement. The c-o-c bond angle (111.7⁰ in methoxy methane) is slightly greater than the tetrahedral angle (109⁰28) due to the repulsive C-O bond length (141 pm) in ethers is almost the same as in alcohols (142 pm in methanol).

pH of Buffer solution

The pH of acidic and basic buffer can be calculated by Henderson – Hasselbalch equations. Consider an acidic buffer HA + A^-
HA <=======> H⁺ + A^-
K_a = [H⁺] [A^-] / [HA]
[H⁺] = K_a [HA]/[A^-]
[H⁺] = K_a [acid]/[salt]
There fore pH = -log[H⁺]
pH = pK_a + log [salt]/[acid]
when, [salt]/[acid] = 1 ,
pH = pK_a
Since pK_a of an acid is a constant at constant temperature, the pH of the buffer is constant. Thus buffer capacity is maximum in a solution containing equivalent amount of acid and its salt.
The pH of basic buffer is also given by Henderson – Hasselbalch equation
BOH <=======> B⁺ + OH^-
K_b = [B⁺][OH^-]/[BOH]
[OH^-] = K_b [BOH]/[B⁺]
pOH = pK_b + log [salt]/[base]
pH = 14 – pOH
= 14 – [pK_b + log [salt]/[base]]

Buffer action

The property of a buffer solution to resist change in its pH value even when small amounts of the acid or the base are added to it is called buffer action.
Consider the acidic buffer solution containing acetic acid and sodium acetate. They dissociate as
CH3COONa <=======> CH3COO_- + H₊
CH3COONa <=======> CH3COO_- + Na₊
When a few drops of an acid, HCl is added to this buffer solution, the H₊ ions combine with CH3COO_- ions to form weakly ionized molecules of CH3COOH.
CH3COO_- + H₊ <=======> CH3COOH
Thus H₊ ion concentration does not change and hence the pH of the solution remains constant.
When a few drops of base, NaOH is added to the buffer solution, hydroxyl ions of the base neutralize the acid, forming salt and water.
Similarly, in a basic buffer solution of NH4OH and NH4Cl, they dissociates as
NH4OH <======> NH4₊ + OH_-
NH4Cl ----------> NH4₊ + Cl_-
When a few drops of a base added, the OH_- ions given by it combine with NH4₊ ions to form the weakly ionized NH4OH.
NH4₊ + OH_- ---------> NH4OH
Thus the OH_- ion concentration or the pH of the solution remains unaffected.
When a small amount of an acid is added, the H₊ ions given by it combines with the OH_- ions already produced by NH4OH.
H₊ + OH_- --------> H2O
Therefore the H₊ ions concentration or the pH of the solution remains unaffected.
The buffer capacity of a buffer solution is defined as the number of moles of acid or base added per liter of the solution to change the pH by one unit.

Buffer solutions

Maintenance of PH in blood and in intracellular fluids is absolutely crucial to the processes that occur in living organisms. This is primarily because the functioning of enzymes is sharply pH dependent. The normal pH value of blood plasma is 7.4 and several illness or death can result from sustained variations of a few tenths of pH unit. Also many medical and cosmetic formulations require that these must be kept and administered at a particular pH. There are solutions which resist the change in pH on addition of small amount of acid or alkali and are called Buffer solution. For example a mixture of H₂CO₃ and HCO₃- is a natural buffer system which maintains the pH of blood. A buffer that is widely used in clinical laboratory and in biochemical studies in the physiological pH range is prepared from tris amino methane (hydroxy methyl) (THAM) [(HOCH₂)3CNH₂].
In order for a solution to act as a buffer it must have two components, one of which is able to neutralize acid and the other able to neutralize the base. Common buffer solutions are mixtures containing a
Weak acid and its conjugate base (one of its salt) called acidic buffer
eg:- CH₃OOH/CH₃COONa, H₂CO₃/Na₂CO₃, Boric acid/borax
Weak base and its conjugate acid (one of its salt) called basic buffer
Eg:- NH4OH/NH4Cl, Zinc hydroxide/ zinc chloride, Glycine/ glycine hydrochloride.

Application of solubility product and common ion effect

Qualitative analysis of cations is largely based on the principle of solubility product and common ion effect. Cations are separated in to six groups depending on the solubility of their salts.
Group-1 as insoluble chlorides
Only Ag⁺, Hg²⁺ and Pb²⁺ form insoluble chlorides since they have low values of K_sp.
Group-2 as insoluble sulphide in acidic medium
H₂S <========> H⁺ + HS^- ;
K₁ - first ionization constant
HS^- <========> H⁺ + S^2-;
K₂ – second ionization constant
[S^2-] = K₁K₂ [H₂S]/[H⁺]²
K_sp values of second group sulphides (PbS, CuS, SnS, HgS, As₂S₃, Bi₂S₃, Sb₂S₃) are very low. In acidic buffer, [S^2-] is decreased due to common ion effect and this results in the precipitation of Pb²⁺, Cu²⁺ etc of second group as their sulphides. Third and fourth group sulphides have high value of K_sp, hence they remain soluble.
Group- 3 as insoluble hydroxide in basic buffer of NH₄OH and NH₄Cl
The concentration of OH^- in ammoniacal solution decreases when NH₄Cl is added to it, because of the common ion effect. Thus only for least soluble hydroxide of Fe³⁺, Al³⁺, Cr³⁺ etc. ionic product exceeds the corresponding solubility products, hence only these ions are precipitated. Hydroxides of successive remain soluble due to high K_sp values.

The solubility product constant (Ksp)

The solubilityof ionic solids in water varies depending on a number of factors like lattice enthalpy of the salt and tha solvation enthalpy of the ions in a solution. As a general rule, for a salt to be able to dissolve in a particular solvent, its solvation enthalpy must be greater than its lattice enthalpy. Each salt has its characteristic solubility, which depends on temperature. We can classify salts on the basis of their solubility in three categories.
Soluble - Solubility > 0.1 M
Slightly soluble - 0.01 M < solubility < 0.1M
Sparingly soluble – solubility < 0.01M
We have now consider the equilibrium between the sparingly soluble ionic salt and its saturated aqueous solution. A solution which remains in contact with excess of the solute is said to be saturated. The amount of a solvent () in 100 ml or 1L) to form a saturated solution at a given temperature is termed as the solubility of the solute in the solvent at that temperature.
For a sparingly soluble salts like AgCl,PbI₂, BaSO₄ etc. Ionisation is very small and concentration of the salt may be considered as constant. Thus, for AgCl,
AgCl <========> Ag⁺ + Cl^-
K = [Ag⁺][Cl^-]/[AgCl]
For a pure solid substance the concentration remains constant and we can write
K_sp = K[AgCl]
=[Ag⁺][Cl^-]
Where K_sp is the solubility product constant or solubility product.
If S represents solubility of AgCl (in mol L^-1)
Then [Ag⁺]= [Cl^-] = S molL^-1
K_sp = [Ag⁺][cl^-] = S²
S = √K_sp
A solubility product constant expression is the product of the chemical equation for a solubility equilibrium, with each term raised to the power given by the coefficient in the chemical equation.
Example:-
For BaSO₄ (binary solute giving two ions)
BaSO₄ <=======> Ba²⁺+ SO₄2^-
Ksp = [Ba²⁺][SO₄2^-] = S2
For PbI₂ (Ternary solute giving three ions)
PbI₂ <=======> Pb²⁺ + 2I^-
Ksp = [Pb²⁺][I^-]2 = (s)(2s)2 = 4S³
For Al(OH)₃ (Quaternary solute giving four ions)
Al(OH)₃ <=======> Al³⁺ + 3OH^-
K_sp = [Al³⁺][OH^-]³ = (S)(3S)³ = 27Ss4
For a solute AxBy, (giving (x+y) ions)
K_sp = [A^y+]^x [B^x-]^y = (xS)^x (yS)^y = X^xY^y(S)^(x+y)
Precipitation reactions
For the sparingly soluble solute AB,
AB <=======> A⁺ +B^-
Q is the equilibrium constant with the given [A⁺] and [B^-]
Q = [A⁺][B^-]
When applied to solubility product, Q is generally called the ionic product. Precipitation occurs if Q>K_sp, then solution is said to be super saturated and precipitation cannot occur if Q_sp (a case of unsaturated solution) and a solution is just saturated if Q=K_sp.

Common ion effect

The decrease in the ionization of a weak electrolyte by the presence of a common-ion from a strong electrolyte, is called the common ion effect. Ionisation of CH₃COOH (weak acid) is decreased by the addition of CH₃COONa (CH₃COO- being the common ion)
CH₃COOH <======> CH₃COO^- + H⁺ ………………………..(A)
CH₃COONa -----------> CH₃COO^- + Na⁺
In the presence of CH₃COO^- equilibrium (A) shifts in backward direction.
Ionisation of H₂S (weak acid) is decreased by the addition of HCl (H⁺ being the common ion)
H₂S <=======> 2H⁺ + s^2-
HCl <=======> H⁺ + Cl^-
Ionisation of NH₄OH (weak base) is decreased by the addition of NH₄Cl (NH4 ⁺ being the common ion)
NH₄OH <=======> NH₄⁺ + OH^-
NH₄Cl --------------> NH₄⁺ + Cl^-
Solubility of a sparingly soluble salt is decreased by the addition of common ion. Presence of AgNO₃ or KCl decreases the solubility of AgCl.
AgCl <=======> Ag⁺ + Cl^-
AgNO₃ <========> Ag⁺ (common ion) + NO₃^-
KCl <=======> K⁺ + Cl^-(common ion)
The common ion effect is thus based on Le-Chatelier’s principle in which the stress on the equilibrium that results from raising one of the product concentrations is relieved by shifting the equilibrium to left.