When 0.243 g of Mg metal is combined with enough HCl to make 100 mL of solution in a constant-pressure calorimeter, the following reaction occurs:

**Reaction: Mg _{(s) }+ 2 HCl_{(aq)} → MgCl_{2(aq)} + H_{2(g)}**

heat released by the reaction = heat absorbed by the solution

**ΔH = -q**

assuming volume solution = mass solution

$\mathbf{q}\mathbf{=}\mathbf{mc}\mathbf{\u2206}\mathbf{T}\phantom{\rule{0ex}{0ex}}\mathbf{q}\mathbf{=}(100g)(4.18\frac{J}{g\xb0C})(34.1\xb0C-23.0\xb0C)\phantom{\rule{0ex}{0ex}}\mathbf{q}\mathbf{=}\mathbf{(}\mathbf{100}\mathbf{}\overline{)\mathbf{g}}\mathbf{)}\mathbf{(}\mathbf{4}\mathbf{.}\mathbf{18}\frac{\mathbf{J}}{\overline{)\mathbf{g}}\overline{)\mathbf{\xb0}\mathbf{C}}}\mathbf{)}\mathbf{(}\mathbf{11}\mathbf{.}\mathbf{1}\overline{)\mathbf{\xb0}\mathbf{C}}\mathbf{)}$

When 0.243 g of Mg metal is combined with enough HCl to make 100 mL of solution in a constant-pressure calorimeter, the following reaction occurs:

Mg_{(s) }+ 2 HCl_{(aq)} → MgCl_{2(aq)} + H_{2(g)}

If the temperature of the solution increases from 23.0 °C to 34.1 °C as a result of this reaction, calculate ΔH in kJ/ mol of Mg. Assume that the solution has a specific heat of 4.18 J/g°C.

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