When 0.250 mol of CH4(g) reacts with excess Cl2(g) at constant pressure according to the chemical equation shown below, 177 kJ of heat are released. Calculate the value of ΔH for this reaction, as written. 2 CH4(g) + 3 Cl2(g) –> 2 CHCl3(l) + 3 H2(g) Group of answer choices-1616 kJ-1716 kJ-1516 kJ-1416 kJ
Question
When 0.250 mol of CH4(g) reacts with excess Cl2(g) at constant pressure according to the chemical equation shown below, 177 kJ of heat are released. Calculate the value of ΔH for this reaction, as written. 2 CH4(g) + 3 Cl2(g) –> 2 CHCl3(l) + 3 H2(g) Group of answer choices-1616 kJ-1716 kJ-1516 kJ-1416 kJ
Solution
The reaction as written involves 2 moles of CH4. However, the problem states that only 0.250 moles of CH4 are used. Therefore, the heat released (ΔH) from 0.250 moles will be a quarter of the heat released from 2 moles.
Given that 177 kJ of heat are released from 0.250 moles, we can calculate the heat released from 2 moles by multiplying 177 kJ by 4 (since 2 moles is four times greater than 0.250 moles).
177 kJ * 4 = 708 kJ
However, the problem asks for the value of ΔH, which is the heat of the reaction. Heat released in an exothermic reaction is represented as a negative value, so ΔH = -708 kJ.
None of the provided answer choices match this calculation. There may be a mistake in the problem or the answer choices.
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