definitions and pH calculations

How is acid strength measured

By hydrogen ion concentration using a negative logarithmic scale

Define pH

pH = -log[H+(aq)]

Define strong acid and give an example

A proton donor which is fully dissociated in water (eg inorganic acids - HNO3, HCl, H2SO4)

Define strong base and give an example

A proton acceptor which is fully dissociated in water (eg hydroxides)

Define weak acid and give an example

A proton donor which is only partially dissociated in water (eg organic acids - carboxylic acids, carbonic acids)

Define weak base and give an example

A proton acceptor which only partially dissociates (eg ammonia, ethanamine)

Explain the strength of a conjugate base in terms of the acid it comes from

• the stronger the acid, the weaker the conjugate base

  • HCl → H+ + Cl-

  • equilibrium wants to shift right, more stable dissociated so weak base Cl- is formed

• The weaker the acid, the stronger the conjugate base

  • CH3COOH ← → H+ + CH3COO-

  • More stable as its acid form, CH3COO- more likely to accept a proton so strong base CH3COO- formed

How much energy is released when a strong acid is neutralised with a strong base and why

• approx -57kJ/mol

• Due to formation of water molecules (H+ + OH-)

How does neutralisation enthalpy differ when a weak acid is neutralised

• some of the energy needed to dissociate the molecules to make the H+ ions available for neutralisation

• Value for neutralisation enthalpy is lower than -57kJ/mol (ie more positive/closer to 0)

What does a change in one pH unit represent in the H+ ion concentration

10-fold change in H+ conc

How is the pH of strong acids found (+why)

Strong acids are completely dissociated so concentration of acid is same as concentration of H+ ions (note: for diprotic, x2 the conc, x3 for triprotic etc)

How is the pH of weak acids found (+why)

• weak acids only partially dissociated so [HA] ≠ [H+]

• HA + H2O → A- + H3O+

• Kc = [A-][H+]/[HA][H2O]

• Ka = [A-][H+]/[HA]

Why is Ka not defined for strong acids

There will be no undissociated acid left

Define pKa

pKa = -log(Ka)

What does a large Ka value mean

The larger the Ka value, the smaller the pKa value therefore the stronger the acid (Larger Ka means higher conc of H+)

Why is pKa/Ka a better measure of acid strength than pH

Their values do not depend on concentration

What is the simplification made when finding the pH of weak acids using Ka

Since 1HA → 1H+ + 1A- every time HA dissociates, at equilibrium, [H+]=[A-]

What are the approximations made when using Ka to find pH of weak acids

• the [H+] due to the dissociation of water is negligible

• The degree of dissociation of a weak acid is negligible

  • [HA]int = [HA]eq

Write the expression for Ka and its approximated version

• Ka = [H+][A-]/[HA]

• Ka = [H+]²/[HA]

What is the assumption made for the ionic product of water

[H2O] is very large and therefore constant, therefore:

Kw = Kc[H2O] = [H+][OH-]

What is Kw

The ionic product constant for water

What is the ionic constant for water at 298K

Kw = 1.0×10^-14 mol²/dm^6

What happens to the acidity and reactivity of water with increased temperature

• Acidity remains unchanged as [H+] = [OH-]

• Increases in reactivity as [H+] increases

What is a simplification for the ionic constant of water for pure water

Kw = [H+]²

What happens to Kw and pH as temperature increases

Kw increases due to higher dissociation of water therefore higher [H+]

How can the pH of strong bases be calculated

using Kw = [H+][OH-] to find [H+]

When diluting a weak acid, is the change in pH less than or more than expected

Less than expected as the acid doesn’t fully dissociated