Question 1

What does the first law of thermodynamics state?

Accepted Answer

The first law states that the change in internal energy of a system equals the heat added to it plus the work done on it: ΔU = q + w. Energy is conserved, so this equation tracks transfers into and out of the selected system.

Question 2

Why is work negative during expansion?

Accepted Answer

During expansion, the system does work on its surroundings rather than receiving work. Under the chemistry sign convention, w = -P_ext ΔV; expansion has positive ΔV and therefore negative w, while compression gives negative ΔV and positive w.

Question 3

When does the heat absorbed at constant pressure equal enthalpy change?

Accepted Answer

The relation ΔH = q_p applies to a constant-pressure process with only pressure-volume work. Under constant volume, ΔV is zero, w is zero, and the corresponding heat relation is ΔU = q_v.

Question 4

Why does Hess's law make enthalpy change path-independent?

Accepted Answer

Enthalpy is a state function, so its change depends only on the initial and final states. A direct route and a multi-step route between the same states have the same total ΔH, and a closed signed cycle has a sum of zero.

Question 5

How are average bond enthalpies used to estimate a reaction enthalpy?

Accepted Answer

Add the enthalpies required to break every bond occurrence in the gaseous reactants, then subtract the enthalpies released when every product bond is formed. The result is an average bond-enthalpy estimate and may differ slightly from an experimental reaction enthalpy.

Energy Ledger — First Law, ΔH vs ΔU & Hess's Law

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