Chemistry: Ionic Compounds: Auto and Self Ionization of Water (For CBSE, ICSE, IAS, NET, NRA 2022)

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Auto-Ionization or Self-Ionization of Water

  • The water has capacity to act as both weak acid weak base.
  • The water molecules present in water undergoes auto ionization as a result a small concentration of H3O + and OH are formed in water.
  • The self-ionization of water can be represented as shown below: H2O + H2O H3O + + OH-
  • The equilibrium constant is given as:

= [H3O + ] [OH-] = Keq x [H2O] 2 = Kw

  • The concentration of water is constant, so Kw, dissociation or ionic product constant of water is used.
  • The value Kw at 298K is measured by its electric conductivity and it is found to be 1.0 x 10-14mol2dm-6

Acidic, Basic and Neutral Solutions

  • In acidic solution hydrogen ion (or hydronium ion) concentration is greater than the hydroxide ion concentration.
  • A basic solution is one in which [OH ] exceeds [H3O + ] and a neutral solution is one in which [OH ] equals [H3O + ] .
  • Neutral solution [H3O + ] = [OH ]
  • Acidic solution [H3O + ] > [OH ]
  • Basic solution [H3O + ] < [OH ]
  • The product [H3O + ] [OH ] is constant and the concentrations of [H3O + ] and [OH ] are not independent but they are linked by the relationship [H3O + ] [OH ] = Kw
  • The self-ionization equilibrium is applicable for pure water and also self-ionization of water in any aqueous solution.
  • The hydronium ions and hydroxide ions are present in all aqueous solution, and they are always in equilibrium with the water molecules.

PH Scale

  • In aqueous solutions the acids and bases concentration of H3O + and OH ions may vary from about 10 M to 10 − 14 M.
  • It is inconvenient to express these concentrations by using powers of 10.
  • So, in 1909 a Danish botanist S. P. L. Sorensen proposed a logarithmic scale or pH scale for expressing the concentrations of H + ions.
PH Scale

The pH is defined as the negative logarithm of the molar concentration of hydrogen ions. pH =- log10 [H + ]

The pH of the neutral solution or pure water is 7 i.e.. [H3O + ] = [OH ] = 1 × 10 – 7 pH = – log 1 x 10 – 7 = 7.0

The pH of the acidic solutions is < 7 i.e.. [H3O + ] > [OH ]

[H3O + ] > 1 × 10 – 7

pH = – log (> 1 × 10 – 7) =< 7.0

The pH of the basic solutions is > 7 i.e.. [H3O + ] < [OH ]

[H3O + ] < 1 × 10 – 7

pH = – log (< 1 × 10 – 7) > 7.0

  • A strongly acidic solution have a pH of less than zero (i.e.. , negative) and a strongly alkaline solution can have a pH value greater than 14.
  • The pH range normally observed is between 0 to 14.
  • The notation p is used in a number of places and it is defined as ‘the negative logarithm of’ .
  • It has been extended to OH (aq) and equilibrium constants like, Ka, Kb and Kw, etc.

pOH =- log10 [OH-]

pKa =- log10 Ka

pKb =- log10 Kb

pKw =- log10 Kw

The Effect of Common Ions on Dissociation of Weak Acids and Bases

  • According to Le Chatlier՚s principle the presence of common ions in a solution of a weak acid or a base will affect by suppressing the dissociation the acid or base.
  • The presence of common ions would make the equilibrium to go to the left and the common ions suppress the equilibrium.

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