In the presence of excess oxygen, methane gas burns in a constant-pressure system to yield carbon dioxide and water: CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (l) with ΔH = -890.0 kJ. Calculate the value of q (kJ) in this exothermic reaction when 1.70 g of CH4 is burned.

• A. -890.0 kJ
• B. -94.5 kJ
• C. -178.0 kJ
• D. -445.0 kJ

Answer: (B)

Verified step by step guidance

1. Determine the molar mass of methane (CH4) by adding the atomic masses of carbon (C) and hydrogen (H). The atomic mass of C is approximately 12.01 g/mol and H is approximately 1.01 g/mol. Therefore, the molar mass of CH4 is 12.01 + (4 * 1.01) g/mol.
2. Calculate the number of moles of CH4 in 1.70 g by using the formula: \( \text{moles of CH4} = \frac{\text{mass of CH4}}{\text{molar mass of CH4}} \).
3. Use the stoichiometry of the reaction to relate the moles of CH4 to the enthalpy change (ΔH). The given ΔH is for the combustion of 1 mole of CH4, which is -890.0 kJ.
4. Calculate the heat released (q) by multiplying the moles of CH4 by the enthalpy change per mole: \( q = \text{moles of CH4} \times (-890.0 \text{ kJ/mol}) \).
5. Since the reaction is exothermic, the value of q will be negative, indicating that heat is released during the reaction.