In a calorimeter of heat capacity $20.0 J K^{- 1}$ , $50.0 mL$ of $1.00 M$ weak acid HA is mixed with $50.0 mL$ of $1.00 M$ NaOH; the temperature rises by $6.00 K$ . Assume the solution has density $1.00 g m L^{- 1}$ and specific heat $4.18 J g^{- 1} K^{- 1}$ . If the enthalpy of neutralisation of a strong acid by a strong base is $- 57.1 kJ mo l^{- 1}$ , the enthalpy of ionisation of HA is

Answer: A

Heat released

= (100 \times 4.18 + 20.0) (6.00) = 2628 J

for

0.0500 mol

of water formed, so

Δ_{n e u t} H = - 52.56 kJ mo l^{- 1}

. This is less negative than

- 57.1

because part of the energy is used to ionise HA:

Δ_{i o n} H = Δ_{n e u t} H (measured) - (- 57.1) = 4.54 kJ mo l^{- 1}