Answer to Question #205388 in General Chemistry for james

Question

If 2.0 moles of H2O(g) are converted to H2(g) and O2(g) against a pressure of 1.0 atm at 125˚C, what is the ∆E for this reaction?

Answer 1

Reaction:

$2H_2O → 2H_{2 (g)} + O_{2 (g) }$

Given Parameters:

H = 483.6 kJ/mol

Number of moles of H2O = 2.0 moles

Pressure (P) = 1.0 atm

Temperature (T) = 125°C

$ΔU = ΔH − RTΔn$

where;

ΔH = change in enthalpy.

R = gas constant = 8.314 J/mol⋅K

T = Temperature

n = number of moles

In the reaction, we see initially there are 2.0 moles of gas. Since products are 2.0 moles of H₂ and 1.0 mole of O₂, there is a net gain of 1 mole of gas (2 reactants → 3 product). Thus, Δn = +1.

Given ΔH = 483.6 kJ, which means it requires 483.6 kJ to decompose 2.0 moles of water

ΔU = 483.6 × 103 J − (8.314 J/mol⋅K398 K+1 mol) = 4.80 × 10⁵ J = 4.80 × 10² kJ

Hence, ΔU for the reaction is 4.80 × 10² kJ.