Question

a chemist measures the energy change \\( \delta h \\) during the following reaction:

\\( \mathrm { cl } _ { 2 } ( g ) + \mathrm { h } _ { 2 } ( g ) \
ightarrow 2 \mathrm { hcl } ( g ) \quad \delta h = - 184. \mathrm { kj } \\)

use the information to answer the following questions.

this reaction is...	\\( \bigcirc \\) endothermic. \\( \bigcirc \\) exothermic.

| \\( \

$$\begin{array} { l } { \\text { if you said heat will be released or absorbed in } } \\\\ { \\text { the second part of this question, calculate how } } \\\\ { \\text { much heat will be released or absorbed. } } \\\\ { \\text { be sure your answer has the correct number of } } \\\\ { \\text { significant digits. } } \\end{array}$$

\\) | \\( \square \mathrm { kj } \\) |

Explanation:

Step1: Determine the type of reaction

When \(\Delta H<0\), the reaction is exothermic. Here \(\Delta H=- 184\space kJ<0\), so the reaction is exothermic.

Step2: Calculate the moles of \(Cl_{2}\)

The molar mass of \(Cl_{2}\) is \(M = 70.90\space g/mol\). The number of moles \(n=\frac{m}{M}\), where \(m = 71.0\space g\). So \(n=\frac{71.0\space g}{70.90\space g/mol}\approx1.00\space mol\)

Step3: Calculate the heat released

From the reaction \(Cl_{2}(g)+H_{2}(g)\to2HCl(g)\space\Delta H=-184\space kJ\), when \(1\space mol\) of \(Cl_{2}\) reacts, \(184\space kJ\) of heat is released. Since \(n(Cl_{2}) = 1.00\space mol\), the heat released \(q = 184\space kJ\)

Answer:

• This reaction is exothermic.
• Yes, released.
• \(184\space kJ\)