what is the relationship between the amount of heat produced per mole of reactant?
The relationship between the amount of heat produced (ΔH) and the number of moles of reactant can be determined by using the concept of molar enthalpy (heat) of reaction.
To find the amount of heat produced per mole of reactant, you can calculate the molar enthalpy of reaction by dividing the change in enthalpy (ΔH) by the number of moles of reactant involved in the balanced chemical equation.
Here's how you can do it step by step:
1. Start by determining the balanced chemical equation for the reaction of interest. This equation will give you the stoichiometry, or the ratio of reactants and products, which is crucial for determining the moles involved.
2. Identify the reactant for which you want to calculate the heat. Make sure you have the number of moles of this reactant.
3. Determine the change in enthalpy (ΔH) for the balanced chemical equation. ΔH is usually given in kilojoules (kJ) or joules (J) per mole of reaction. It represents the heat absorbed or released during the reaction.
4. Calculate the molar enthalpy of reaction by dividing the change in enthalpy (ΔH) by the number of moles of the reactant involved.
For example, let's consider the reaction of methane combustion:
CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g)
Suppose we want to find the heat produced per mole of methane (CH₄).
1. The balanced chemical equation tells us that 1 mole of CH₄ reacts with 2 moles of O₂.
2. Let's assume we have 0.5 moles of CH₄.
3. The change in enthalpy (ΔH) for methane combustion is -890.3 kJ/mol (a negative value indicates heat release).
4. To find the heat produced per mole of CH₄, divide the change in enthalpy (ΔH) by the number of moles of CH₄:
Heat produced per mole of CH₄ = ΔH / moles of CH₄
= -890.3 kJ/mol / 0.5 mol
= -1780.6 kJ/mol
So, in this example, the heat produced per mole of CH₄ is -1780.6 kJ/mol.
Remember to pay attention to the sign convention. A negative value indicates heat released, while a positive value represents heat absorbed.