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A

Anonymous

Jan 22, 2025

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

Lighters are typically 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. What are the values of ΔH and ΔE for the combustion of one mole of butane? Could you please explain how to calculate these values?

1 Answers

A
Anonymous

Feb 02, 2025

So, consider the necessary formulas:

deltaH = q - PdeltaV

deltaE = q + w

First, to find delta E:

The reaction PRODUCES 2658 kJ of h (q), and does 3 kJ of work (w).

2658 kJ(q) + 3 kJ(w) = 2661 kJ, BUT the reaction PRODUCES heat, which means deltaE is negative.

deltaE = -2661

Second, to find deltaH:

deltaH = q - PdeltaV

deltaH = 2658 kJ(q) - PdeltaV

Now, the question states that butane burns at a constant pressure; that just translates to the pressure of the reaction is equal to 0.

deltaH = 2658 kJ(q) - (0)deltaV

deltaH = 2658 kJ(q) - 0

deltaH = 2658 kJ, BUT, like before, the reaction PRODUCES heat, which also mean deltaH is negative.

deltaH = -2658 kJ

I hope this helped! Have a nice week.

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