The pH scale is a logarithmic scale that is used to measure the acidity or the basicity of a substance. The possible values on the pH scale range from 0 to 14. The term ‘pH’ is an abbreviation of ‘potential for hydrogen’ or ‘power of hydrogen’. Acidic substances have pH values ranging from 1 to 7 (1 being the most acidic point on the pH scale) and alkaline or basic substances have pH values ranging from 7 to 14. A perfectly neutral substance would have a pH of exactly 7.
The pH of a substance can be expressed as the negative logarithm (with base 10) of the hydrogen ion concentration in that substance. Similarly, the pOH of a substance is the negative logarithm of the hydroxide ion concentration in the substance. These quantities can be expressed via the following formulae:
- pH = -log[H+]
- pOH = -log[OH–]
It is important to note that the pH scale is a logarithmic scale. Therefore, an increase in the pH value of a solution by one will be accompanied by a tenfold increase in the hydrogen ion concentration and therefore, a tenfold increase in acidity. Elaborating with the help of an example, a solution with a pH of 3 will have ten times the acidity of a solution with a pH of 4 and a hundred times the acidity of a solution with a pH of 5.
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
- 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).
- pH value is zero for 1N solution of strong acid. Concentration of 2 N, 3 N, 10 N, etc. gives negative pH values..
- 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
- 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.
- 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.
- 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).
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
Molarity of water
Molarity = Number of moles per litre of solution = = 55.55 mole l-1
Ionization constant of water
Ka = Kb =
Where, Ka is the acid ionization constant and Kb is the base ionization constant.
pKa = pKb = – log[Ka] =
Degree of ionization of water
Initial concentration moles 55.55 0 0
At equilibrium moles 10-7 10-7
Degree of ionization = α =
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
Also Read: Study The pH Change