Give definitions:

i. Arrhenius acid

ii. Arrhenius base

iii. Bronsted-Lowry acid

iv. Bronsted-Lowry base

v. Concentration of solution

vi. pH of solution

vii. pOH of solution

viii. Strong acid

ix. Weak acid

x. Strong base

xi. Weak base

xii. Neutralisation reaction

xiii. 1M concentration

i. Arrhenius acid is a hydrogen-containing substance which produces hydrogen ion (H ⁺ ) in its aqueous solution.

Examples of such acids are Nitric acid and hydrochloric acid.

HNO₃→ H ⁺ + NO₃^-

HCl → H ⁺ + CL^-

Both contain Hydrogenand produce H ⁺

ii. Arrhenius base is hydroxide containing substance and which produces OH- ion in its aqueous solution.

Examples of two such bases are sodium Hydroxide and potassium hydroxide.

NaOH → Na ⁺ + OH^-

KOH → K ⁺ + OH^-

Both contain Hydroxide and produces OH^-

iii. The substances which donate protons [H ⁺ ] to the other substance are called bronsted-lowry acids. The species responsible for acidity is not [H ⁺ ] but [H₃O ⁺ ].

Examples of bronsted-lowry acids are: HCl, HNO₃

iv. Bronsted-lowry base is a substance which accepts proton from other substance. For example ammonia.

When ammonia is dissolved in the water, it accepts a proton from water and thus behaves as a bronsted-lowry base.

v. A system resulting when a solute is dissolved in a solvent is called a solution. The substance which is present in lesser proportion is called solute and the substance which is in present in a larger proportion is known as a solvent.

The amount of solute in relation to the amount of solvent is

called concentration. The concentration of a solution can be expressed in terms of molarity, normality etc.

vi. The pH is defined as the negative logarithm to the base 10 of the concentration of H₃O ⁺ in aqueous solution. In simple words, pH is used to express the hydronium ion concentration in a solution.

Less the pH more will the hydronium ion concentration.

pH = -log₁₀[H₃O ⁺ ]

vii. The pOH is defined as the negative logarithm to the base 10 of the concentration of hydroxide ions in the aqueous solution. In simple words, pOH is used to express the hydroxide ion concentration in a solution. Less the pOH more will be hydroxide ion concentration.

pOH = -log₁₀[OH^-]

viii. Strong acids are those acids which completely ionize or dissociate in their aqueous solutions.

Examples:

•Hydrochloric acid

•Nitric acid

Dissociation of nitric acid: HNO₃→ H ⁺ + NO₃^-

Dissociation of hydrochloric acid: HCl → H ⁺ + CL^-

ix. Weak acids are those acids which partially ionize or dissociate in their aqueous solutions.

Examples:

•Oxalic acid

•Citric acid

x. Strong bases are those bases which completely ionise or dissociate in their aqueous solutions.

Examples:

•Sodium hydroxide

•Potassium hydroxide.

Dissociation of Sodium hydroxide: NaOH → Na ⁺ + OH^-

Dissociation of potassium hydroxide: KOH → K ⁺ + OH^-

xi. Weak bases are those bases which partially ionize or dissociate in their aqueous solutions.

Examples:

•Ammonium hydroxide

•Ammonia

xii. When a reaction takes place between an acid and a base, salt and water are produced. This reaction is called a neutralisation reaction. In this reaction acid and base mutually nullify each other’s properties. An example of such a reaction is:

xiii. The amount of solute in relation to the amount of solvent is

called concentration. The concentration of a solution can be expressed in terms of molarity, normality etc. 1M concentration can be defined as 1 mole of a substance dissolved in 1 litre of a solution.