A-Level Chemistry OCR Notes
5.1.3 Acids, bases, and buffers
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Brønsted–Lowry Acid and Base
- Brønsted-Lowry Acid is a species that can donate protons.
HA (aq) + H2O (l) → H3O+ (aq) + A- (aq)
- Brønsted-Lowry Base is a species that can accept protons.
B (aq) + H2O (l) → BH+ (aq) + OH- (aq)
- Conjugate acid-base pairs are a pair of species that interconvert by the gain or loss of a proton.
- Monobasic Acids – each molecule can release one proton, e.g. HCl
HCl (aq) → H+ (aq) + Cl- (aq)
- Dibasic Acids – each molecule can release two protons in two steps, e.g. H2SO4
H2SO4 (aq) → H+ (aq) + HSO4- (aq)
HSO4- (aq) → H+ (aq) + SO4^2- (aq)
HSO4- (aq) → H+ (aq) + SO4^2- (aq)
- Neutralisation reactions:
acid + base → salt + water
- Acid-carbonate reactions:
acid + carbonate → salt + carbon dioxide + water
- Acid-metal reactions:
acid + metal→ salt + hydrogen gas
pH
- The concentration of H+ ions, often written as [H+], present in a solution determines how acidic it is.
- The pH scale is logarithmic to cover the wide range of hydrogen ion concentrations in aqueous solution
- pH = - log10 [H+]
- The greater the concentration of H+ ions, the lower the pH value, and the more acidic the solution is
- [H+] = 10 -pH
- The pH of a strong acid can be calculated from its concentration:
- Strong acids fully dissociate: HA (aq) → A- (aq) + H+ (aq)
- Therefore, the concentration of H+ ions is equal to the concentration of the acid before dissociation: [HA] = [H+]
- This is only the case for monoprotic acids, such as HBr. Diprotic acids, such as H2SO4, will require the concentration of the acid before dissociation to be multiplied by two to get the concentration of H+ ions.
Ionic Product of Water
- Species that can act as both an acid and a base are called amphoteric, e.g. water
- Water slightly dissociates into hydroxide and hydroxonium ions.
2H2O (I) ⇌ OH- (aq) + H3O- (aq)
H2O (I) ⇌ OH- (aq) + H+ (aq)
H2O (I) ⇌ OH- (aq) + H+ (aq)
- The equilibrium constant for the dissociation of water is:
The concentration of H2O is constant for a given temperature
- Kw, like other equilibrium constants, only changes with temperature.
- Kw can be used to calculate the pH of a strong base from its concentration
- Strong bases dissociate almost fully, meaning the concentration of OH- ions is equal to the concentration of the base before dissociation:
Buffers
- A buffer is a solution that minimises changes in pH upon dilution or on the addition of a small amount of an acid or a base.
- Acidic buffers are a mixture of a weak acid and the salt of the same weak acid
- Basic buffers are a mixture of a weak acid and a strong alkali
- The pH of a buffer solution can be calculated if the concentration of H+ ions is known:
- It Is assumed the concentration of the acid before dissociation is equal to the concentration of the acid at equilibrium:
- It is also assumed the salt is ionic so will fully dissociate:
- The carbonic acid-hydrogencarbonate ion buffer is an important buffer that regulates the pH of blood
Strong and Weak Acids
- A strong acid is an acid which dissociates almost completely in water or aqueous solution.
- HA (aq) → A- (aq) + H+ (aq)
- A weak acid is an acid which is only partially dissociated in water or aqueous solution.
- HA (aq) ⇌ A- (aq) + H+ (aq)
- The dissociation constant for a weak acid, Ka, is used to measure the strength of a weak acid:
[A-] conjugate base concentration (mol dm^-3)
[H+] H+ ion concentration (mol dm^-3)
[HA] acid concentration (mol dm^-3)
[H+] H+ ion concentration (mol dm^-3)
[HA] acid concentration (mol dm^-3)
- The greater the strength of the weak acid, the greater its Ka value.
- The pH of a weak acid can be calculated using the concentration of the acids and Ka
- It is assumed [H+] = [A-]
- The concentration of HA present is given by the concentration of HA before any dissociation minus the concentration of H+ ions:
- As the degree of dissociation of the weak acid is very small:
the H+ ion concentration can be used to find the pH
- A logarithmic scale for pH can be used pKa = - log10 Ka
- Ka = 10 -pKa
- The weaker the weak acid, the smaller the Ka value, and therefore, the higher the pKa
- The assumption that [H+] = [A-] fails for high pH conditions, because the dissociation of water is significant, resulting in an increased [H+]
- The calculations also assume that at equilibrium, [HA] >> [H+], which may not be true for stronger weak acids
Titrations
- Acid-base titrations can be used to find the concentration of a sample of either an acid or a base.
- A known concentration of an acid is gradually added to a known volume of a base of unknown concentration until the solution is neutralised
- A pH meter or an indicator can be used to monitor the pH
- The equivalence point is the point at which all of the acid in the known volume has reacted with the base
- Titration curves: a diprotic acid indicates two equivalence points
- Indicators are used to visualise the equivalence point more easily
- An acid-base indicator is a weak acid, HIn, with the equilibrium
HIn (aq) ⇌ In- (aq) + H+ (aq)
- The indicator and its conjugate base are different colours in solution. Most indicators change colour over a range of two pH units spread around their end point
- An indicator must have:
- A sharp colour change
- An end point close to the equivalence point
- A clear colour change
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