Question

A student investigated the enthalpy of combustion (deltaHc) of methanol under standard conditions using the apparatus shown in the diagram. The measurements the student recorded are shown in the table. Use this information to answer the questions below.

_Clamp stand
_
_ _I__Thermometer
_ I I Beaker
_ I-----I
_ I___I Water
_ I Wick
_ -----
_ I I
_ ----- Alcohol
_ I __I
______________

Alcohol mass before burning: 80.6g
Alcohol mass after burning: 75.9g
Water heated: 100g
Methanol Mr: 32
Initial temperature of water: 21.5 C
Final temperature of water: 32.4 C
Enthalpy of combustion of 1 mole of methanol:

The student wanted to know if the value obtained in part 1 is similar to that calculated using average bond enthalpy data.

a) Using the balanced equation and the data in the table below, calculate the theoretical enthalpy of combustion.

Note: you will need to include the enthalpy of vaporisation for the liquid components which are also given.

CH3OH(l) + 1.5O2(g) → CO2(g) + 2H2O(l)

Average Bond Enthalpies KJmol-1
C-H 412
C-C 348
C-O 358
O=O 496
C=O 743
O-H 463

Enthalpy of vaporisation KJmol-1
Methanol 35
Water 41

b) Suggest some reasons as to why this value is different to the one obtained in the practical. (What are the reasons that the the theoretical and actual enthalpy changes are different)

Answer 1

The Molar enthalpy of combustion of methanol obtained from the practical is -31.02 kJ/mol.

The molar enthalpy of combustion from bond energies, ΔH is -543 kJ/mol

The differences in the values could be due to:

1. Heat losses from the calorimeter to the surroundings
2. Inaccurate measurements of mass and temperature changes

What is the enthalpy theoretical and actual enthalpy of combustion of methanol?

The actual enthalpy of combustion of methanol is calculated from the data obtained from the laboratory work.

The enthalpy of combustion is equal to the heat energy given off from the combustion of methanol.

Quantity of heat gained by water , q = -mcθ

where:

• m = mass
• c = specific heat capacity
• θ = temperature change

mass of water heated = 100 g

specific heat capacity of water = 4.18 J/g/°C

temperature change = 32.4°C - 21.5°C = 10.9°C

q = -(100 * 4.18 * 10.9) J

q = - 4556.2 J

moles of methanol reacted = mass/molar mass

mass of methanol = 80.6 - 75.9 = 4.7 g

molar mass of methanol = 32 g/mol

moles of methanol = 4.7/32 = 0.14687 moles

Molar enthalpy of combustion of methanol = -4556.2 J/0.14687 mole

Molar enthalpy of combustion of methanol = -31.02 kJ/mol

Calculating molar enthalpy of combustion from bond energies:

• ΔH = sum of the bond energies of bonds broken - sum of the bond energies of the bonds formed

Sum of bond energies of bonds broken = (3 * 412) + 358 + 463 + (1.5 * 496) + 35 = 2836 kJ/mol

Sum of bond energies of bonds formed = (2 * 743) + 2(2 * 463) + 41 = 3379 kJ/mol

ΔH = 2836 kJ/mol - 3379 kJ/mol

ΔH = -543 kJ/mol

The value obtained from the practical is less than that from the bond energies

b. The possible reasons for the difference in value obtained from the practical is less than that from the bond energies include:

• Heat losses from the calorimeter to the surroundings
• Inaccurate measurements of mass and temperature changes.

In conclusion, the molar enthalpy of combustion of methanol is negative since heat is given off during the reaction.

Learn more about enthalpy and bond energies at: https://brainly.com/question/13981382

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