Lighters are usually fueled by butane (C4H10). When 1 mole of butane burns at constant pressure, it produces 2658 kJ of heat and does 3 kJ of work.

Part A

What are the values of ΔH and ΔE for the combustion of one mole of butane?

Part B

Express your answer using four significant figures.

Question

Branson Manning

Chemistry

29 March, 00:21

Lighters are usually fueled by butane (C4H10). When 1 mole of butane burns at constant pressure, it produces 2658 kJ of heat and does 3 kJ of work.

Part A

What are the values of ΔH and ΔE for the combustion of one mole of butane?

Part B

Express your answer using four significant figures.

+4

Answers (1)

Know the Answer?

Bucky · Answer 1

see explanation below

Explanation:

This is actually something that you can solve using the following expression:

ΔH = q + PΔV

ΔE = q + W

As the exercise states that the butane is burned at constant pressure, we can assume that P = 0, therefore:

ΔH = q

According to this, ΔH = q = 2658 kJ. However, as it's produced in the reaction, this value is negative, so ΔH = - 2658 kJ

To get ΔE:

ΔE = 2658 + 3 = 2661 kJ

But as before, this is produced in the reaction, so this value is also negative

ΔE = - 2661 kJ

Lighters are usually fueled by butane (C4H10). When 1 mole of butane burns at constant pressure, it produces 2658 kJ of heat and does 3 kJ of work. Part AWhat are the values of ΔH and ΔE for the combustion of one mole of butane?Part BExpress your answer using four significant figures.

Lighters are usually fueled by butane (C4H10). When 1 mole of butane burns at constant pressure, it produces 2658 kJ of heat and does 3 kJ of work.

Part A

What are the values of ΔH and ΔE for the combustion of one mole of butane?

Part B

Express your answer using four significant figures.