Ka, pH & Kw (OCR A Level Chemistry)

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The Acid Dissociation Constant, Ka

Weak acids

A weak acid is an acid that partially (or incompletely) dissociates in aqueous solutions

  • Eg. all organic acids e.g. ethanoic acid and some inorganic acids HCN (hydrocyanic acid), H2S (hydrogen sulfide) and H2CO3 (carbonic acid)
  • The position of the equilibrium is more over to the left and an equilibrium is established

Equilibria Dissociation of a Weak Acid, downloadable AS & A Level Chemistry revision notes

The diagram shows the partial dissociation of a weak acid in aqueous solution

  • As this is an equilibrium we can write an equilibrium constant expression for the reaction

The acid dissociation constant, downloadable AS & A Level Chemistry revision notes

  • This constant is called the acid dissociation constant, Ka, and has the units mol dm-3
  • Values of Ka are very small, for example for ethanoic acid Ka = 1.74 x 10-5 mol dm-3 
  • When writing the equilibrium expression for weak acids, the following assumptions are made:
    • The concentration of hydrogen ions due to the ionisation of water is negligible

  • The value of Ka indicates the extent of dissociation
    • The higher the value of Ka the more dissociated the acid and the stronger it is
    • The lower the value of Ka the weaker the acid

pKa

  • The range of values of Ka is very large and for weak acids, the values themselves are very small numbers

Table of Ka values

Table of Ka values, downloadable AS & A Level Chemistry revision notes

  • For this reason it is easier to work with another term called pKa
  • The pKa  is the negative log of the Ka value, so the concept is analogous to converting [H+] into pH values

pKa = -logKa

  • Looking at the pKa values for the same acids:

Table of pKvalues

Table of pKa values, downloadable AS & A Level Chemistry revision notes

  • The range of pKa values for most weak acids lies between 3 and 7

pH & The Ionic Product of Water, Kw

pH

  • The acidity of an aqueous solution depends on the number of H+ (H3O+) ions in solution
  • The pH is defined as:

pH = -log[H+]

    • where [H+] is the concentration of hydrogen ions in mol dm–3

  • Similarly, the concentration of H+ of a solution can be calculated if the pH is known by rearranging the above equation to:

[H+] = 10-pH

  • The pH scale is a logarithmic scale with base 10
  • This means that each value is 10 times the value below it. For example, pH 5 is 10 times more acidic than pH 6.
  • pH values are usually given to 2 decimal places
  • The relationship between concentration is easily seen on the following table

pH & [H+] Table

pH and [H+] Table, downloadable IB Chemistry revision notes

The ionic product of water, Kw

  • In all aqueous solutions, an equilibrium exists in water where a few water molecules dissociate into protons and hydroxide ions
  • We can derive an equilibrium constant for the reaction:

Deriving Kw, downloadable AS & A Level Chemistry revision notes

  • This is a specific equilibrium constant called the ionic product for water
  • The product of the two ion concentrations is always 1 x 10-14 moldm-6
  • This makes it straightforward to see the relationship between the two concentrations and the nature of the solution:

[H+] & [OH] Table

[H+] and [OH-] table, downloadable IB Chemistry revision notes

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Philippa

Author: Philippa

Expertise: Chemistry

Philippa has worked as a GCSE and A level chemistry teacher and tutor for over thirteen years. She studied chemistry and sport science at Loughborough University graduating in 2007 having also completed her PGCE in science. Throughout her time as a teacher she was incharge of a boarding house for five years and coached many teams in a variety of sports. When not producing resources with the chemistry team, Philippa enjoys being active outside with her young family and is a very keen gardener.