Question

In the presence of excess oxygen, methane gas burns in a constant-pressure system to yield carbon dioxide and water: CH₄(g) + 2 O₂(g) → CO₂(g) + 2 H₂O(l) ΔH = -890.0 kJ. Calculate the value of q (kJ) in this exothermic reaction when 1.90 g of methane is combusted at constant pressure?

Answer 1

To calculate the heat released when 1.90 g of methane is burned, convert the mass of methane to moles and multiply by the ΔH per mole. The result is approximately -105.4 kJ of heat released during the combustion.

Step-by-step explanation:

When methane gas, CH₄, is combusted in excess oxygen, the reaction is highly exothermic, with a known ΔH of -890.0 kJ per mole of methane burned. The question asks to calculate the heat (q) released when 1.90 g of methane is burned at constant pressure. First, we must convert the mass of methane to moles using its molar mass (16.04 g/mol). This calculation is as follows:

• Number of moles of CH₄ = mass (g) / molar mass (g/mol) = 1.90 g / 16.04 g/mol ≈ 0.1184 mol

Since the thermochemical equation indicates that 1 mol of CH₄ releases 890.0 kJ, the heat released by 0.1184 mol can be calculated by proportional scaling:

• Heat released (q) = moles of CH₄ × ΔH = 0.1184 mol × -890.0 kJ/mol ≈ -105.4 kJ

Therefore, when 1.90 g of methane is combusted, approximately -105.4 kJ of heat is released.