Question

Find ΔH for the combustion of methanol (CH4O) to carbon dioxide and liquid water from the following data. The heat capacity of the bomb calorimeter is 34.65 kJ/Kand the combustion of 1.761 g of methanol raises the temperature of the calorimeter from 294.40 K to 295.55 K .

Answer 1

I have compGiven,

The heat capacity of teh calorimeter, c=34.65 kJ/K

The mass of methanol, m=1.761 g=1.761×0.0312=0.055 moles

The rise in the temperature, ΔT=295.55-294.40=1.15 K

The heat lost by the combustion of the methanol is gained by the calorimeter.

Thus the heat gained by the calorimeter is given by,

`\Delta E=c*\Delta T`

On substituting the known values,

`\begin{gathered} \Delta E=34.56*10^3*1.15 \\ =39.74\text{ kJ} \end{gathered}`

Thus the ΔH of the reaction is given by,

`\Delta H=(\Delta E)/(m)`

On substituting the known values,

`\begin{gathered} \Delta H=(39.74*10^3)/(0.055) \\ =722.55(kJ)/(mol) \end{gathered}`

Thus the ΔH is 722.55 kJ/mol