pH Scale and Acidity

Table of Content:

• Chemical Equilibrium
• Ionic Equilibrium – Degree of Ionization and Dissociation
• Equilibrium Constant – Characteristics and Applications
• Le Chatelier’s Principle on Equilibrium
• Solubility and Solubility Product
• Acid and Base
• pH Scale and Acidity
• pH and Solutions
• Hydrolysis, Salts, and Types
• Buffer Solutions

Introduction to pH Scale

Acid solutions have protons and basic solutions have hydroxide ions.

Concentrations of the ions are low (negative power of ten). pH scale is a convenient way of expressing these low concentrations in simple numbers between 1 and 14.

pH is the negative logarithm to the base ten of hydrogen ion concentration in moles per litre.

pH = – log [H⁺]

p(OH) is the negative logarithm to the base ten of hydroxide ion concentration in moles per litre.

p(OH) = – log [OH^–]

In aqueous solutions, pH + p(OH) = 14.

pH scale is based on neutral water, where [H⁺] = [OH^–] = 10^-7

For a neutral solution pH = = – log [H⁺] = – log [10^-7] = +7

pH of strong acid decreases with a limit of 1 and pH of a base increases up to 14

Generally, acids and bases will have a pH between.

But negative and greater than14 pH values are also possible.

Limitations of pH Scale

1. pH values does not reflect directly the relative strength of acid or bases.

A solution of pH = 1 has a hydrogen ion concentration 100 times that of a solution of pH = 3 (not three times). A 4 x 10^-5 N HCI is twice concentrated of a 2 x 10^-5 N HCI solution, but the pH values of these solutions are 4.40 and 4.70 (not double).

1. pH value is zero for 1N solution of strong acid. Concentration of 2 N, 3 N, 10 N, etc. gives negative pH values..
2. A solution of an acid having very low concentration, say 10^-8 N, shows a pH 8and hence should be basic, but actual pH value is less than 7.

Periodic Variation of Acidic and Basic Properties

(a)  Hydracids of the Elements of the Same Periods

Along the period acidic strength increases.. Hydrides become increasingly acidic from CH₄, NH₃, H₂O and HF. The increase in acidic properties is due to the fact that the stability of their conjugate bases increases in the order

CH^–₃< NH^–₂  < OH^– < F^–

(b)  Hydracids of the Elements of the Same Group

1. Acidic nature increases down the column. Hydrides of V group elements (NH₃, PH₃, AsH₃, SbH₃, BiH₃) show basic character which decreases due to increase in size and decrease in electronegativity from N to Bi. There is a decrease in electron density in, sp³ -hybrid orbital and thus electron donor capacity decreases.
2. Hydracids of VI group elements (H₂0, H₂S, H₂Se, H₂Te) act as weak acids. The strength increases in the order H₂0 < H₂S < H₂Se < H₂Te. The increasing acidic properties reflects decreasing trend in the electron donor capacity of OH^–, HS^–, HSe^– or HTe^– ions.
3. Hydracids of VII group elements (HF, HCI, HBr, HI) show acidic properties which increase from HF to HI. This is explained by the fact that bond energies decrease.

(H-F = 135 kcal/mol, HCI = 103, HBr = 88 and HI = 71 kcal/mol).

Oxyacids

The acidic properties of oxyacids of the same element which is in different oxidation states increase with an increase in oxidation number.

+ 1            +3          +5            +7

HCIO  <  HC1O₂   <   HC1O₃   <  HCIO₄

+4            +6         +3            +5

H₂SO₃  <   H₂SO₄;    HNO₂     <  HNO₃

But this rule fails in oxyacids of phosphorus.

H₃PO₂  >  H₃PO₃  >  H₃PO₄

The acidic properties of the oxyacids of different elements which are in the same oxidation state decreases as the atomic number increases. This is due to increase in size and decrease in electronegativity.

HC1O₄ > HBrO₄ > HIO₄

H₂SO₃  > H₂SeO₃

But there are a number of acid-base reactions in which no proton transfer takes place, e.g.,

SO₂ + SO₂  ↔  SO²⁺ +  S

Acid₁  Base₂    Acid₂   Base₁

Thus, the protonic definition cannot be used to explain the reactions occurring in non-protonic solvents such as COCl₂, S0₂, N₂O₄, etc.

Water – Amphoteric Weak Electrolyte

1) Water can behave like acid or a base. So it is amphoteric.

Water accepts proton from HCl and acts as a base.

Water gives proton to ammonia and can be an acid \(NH_3(aq)+H_2O(l)\rightleftharpoons NH_{4}^{+}(aq)+OH^{-}\)

Molarity of water

Molarity = Number of moles per litre of solution = = 55.55 mole l^-1

Ionization constant of water

Ka = Kb = \(\frac{[H+[OH-]]}{[H_2O]}\) = \(\frac{10^{-7}\times 10^{-7}}{55.55}=1.8\times10^{-16}\)

Where, Ka is the acid ionization constant and Kb is the base ionization constant.

pKa = pKb = – log[Ka] = \(-\log 1.8\times10^{-16}\) = 15.74

Degree of ionization of water

^{Initial concentration moles 55.55 0 0}

^{At equilibrium moles} 10^-7 10^-7

Degree of ionization = α =\(\frac{number\;of\;moles\;ionized}{initial\;number\;of\;moles}\) \(=\frac{10^{-7}}{55.55}=1.8\times 10^{-9}\)

Only about 2 parts per billion (ppb) of the water molecules dissociate into ions at room temperature. 

Ionic product of water

It is the product of the concentrations of hydrogen and hydroxide ions in water.

Ionic product of water = Kw = [H⁺][OH^–] = 10^-14

pKw = – log[Kw] = – log10^-14= 14

Ionic product, pKw, pKa and pKb remains the same whether the solution is acidic, neutral or

basic.