Question

A sample of solid pyrene (C16H10) that weighs 0.5063 g is burned in an excess of oxygen to CO2(g) and H2O() in a constant-volume calorimeter at 25.00 °C. The temperature rise is observed to be 2.130 °C. The heat capacity of the calorimeter and its contents is known to be 9.233×103 J K-1.

(a) Write and balance the chemical equation for the combustion reaction. Use the lowest possible coefficients. Use the pull-down boxes to specify states such as (aq) or (s).

(b) Assuming that H° is approximately equal to E, calculate the standard enthalpy change for the combustion of 1.000 mol of pyrene to CO2(g) and H2O(). kJ mol-1

(c) Calculate the standard enthalpy of formation per mole of pyrene, using the following for the standard enthalpies of formation of CO2(g) and H2O(). Hf° H2O () = -285.83 kJ mol-1 ; Hf° CO2(g) = -393.51 kJ mol-1 kJ mol-1

Answer 1

a) Balanced equation:

2C₁₆H₁₀ (s) + 21O₂ (g) --------> 32CO₂ (g) + 10H_{​​​​​​2}O (l)

For 1 mol of C₁₆H₁₀:

C₁₆H₁₀ (s) + (21/2)O₂ (g) --------> 16CO₂ (g) + 5H_{​​​​​​2}O (l)

b)

Heat capacity of calorimeter (C) = 9.233 x 10³ J/K

= 9.233 kJ/K

Change in temperature (∆T) = 2.130 K

Heat released = -Heat absorbed by calorimeter

= -C*∆T

= -(9.233 kJ/K)*(2.130 K)

= -19.67 kJ

Molar mass of pyrene = 202.25 g/mol

Moles of pyrene = 0.5063 g/202.25 g/mol =

Moles of pyrene = 0.002503 mol

Change in enthalpy (∆H_reaction) = -19.67 kJ/0.002503 mol

= -7859 kJ/mol

c) Using standard enthalpies of formation and balanced equation for 1 mole:

∆H_reaction = 16∆H_f(CO₂) + 5∆H_f(H₂O) - ∆H_f(C₁₆H₁₀)

-7859 = 16(-393.51) + 5(-285.83) - ∆H_f(C₁₆H₁₀)

∆H_f(C₁₆H₁₀) = +133.69 kJ/mol