PH

pH is a measure of the activity of hydrogen ions (H+) in a solution and, therefore, its acidity or alkalinity. In aqueous systems, the hydrogen ion activity is dictated by the dissociation constant of water (Kw) = 1.011 × 10−14 at 25 °C) and interactions with other ions in solution. Due to this dissociation constant a neutral solution (hydrogen ion activity equals hydroxide ion activity) has a pH of approximately 7. Aqueous solutions with pH values lower than 7 are considered acidic, while pH values higher than 7 are considered alkaline.

Calculation of pH for weak and strong acids

Values of pH for weak and strong acids can be approximated using certain assumptions.

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Under the Brønsted-Lowry theory, stronger or weaker acids are a relative concept. But here we define a strong acid as a species which is a much stronger acid than the hydronium (H3O+) ion. In that case the dissociation reaction (strictly HX+H2O↔H3O++X− but simplified as HX↔H++X−) goes to completion, i.e. no unreacted acid remains in solution. Dissolving the strong acid HCl in water can therefore be expressed:

Related Topics:
Brønsted-Lowry theory - HCl

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:HCl(aq) → H+ + Cl−

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This means that in a 0.01 mol/L solution of HCl it is approximated that there is a concentration of 0.01 mol/L dissolved hydrogen ions. From above, the pH is: pH = −log10 :

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:pH = −log (0.01)

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which equals 2.

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For weak acids, the dissociation reaction does not go to completion. An equilibrium is reached between the hydrogen ions and the conjugate base. The following shows the equilibrium reaction between methanoic acid and its ions:

Related Topics:
Equilibrium - Conjugate base - Methanoic acid

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:HCOOH(aq) ↔ H+ + HCOO−

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It is necessary to know the value of the equilibrium constant of the reaction for each acid in order to calculate its pH. In the context of pH, this is termed the acidity constant of the acid but is worked out in the same way (see chemical equilibrium):

Related Topics:
Equilibrium constant - Acidity constant - Chemical equilibrium

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:Ka = /

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For HCOOH, Ka = 1.6 × 10−4 (some other Ka values)

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When calculating the pH of a weak acid, it is usually assumed that the water does not provide any hydrogen ions. This simplifies the calculation, and the concentration provided by water, 1×10−7 mol, is usually insignificant.

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With a 0.1 mol/L solution of methanoic acid (HCOOH), the acidity constant is equal to:

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:Ka = /

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Given that an unknown amount of the acid has dissociated, will be reduced by this amount, while and will each be increased by this amount. Therefore, may be replaced by 0.1 − x, and and may each be replaced by x, giving us the following equation:

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:1.6 imes 10^{-4} = rac{x^2}{0.1-x}

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Solving this for x yields 3.9×10−3, which is the concentration of hydrogen ions after dissociation. Therefore the pH is −log(3.9×10−3), or about 2.4.

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More on pH calculation...

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