Question

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

\\( 2 \mathrm{fe}_{2} \mathrm{o}_{3}(s) \
ightarrow 4 \mathrm{feo}(s)+\mathrm{o}_{2}(g) \quad \delta h=560 . \mathrm{kj} \\)

use the information to answer the following questions.

this reaction is	\\( \bigcirc \\) endothermic
suppose \\( 90.5 \mathrm{~g} \\) of \\( \mathrm{fe}_{2} \mathrm{o}_{3} \\) react	\\( \bigcirc \\) yes, absorbed
will any heat be released or absorbed?	\\( \bigcirc \\) yes, released
	\\( \bigcirc \\) no.
if you said heat will be released or absorbed in
the second part of this question, calculate how	\\( \square \mathrm{kj} \\)
much heat will be released or absorbed.
be sure your answer has the correct number of
significant digits.

Explanation:

Step1: Determine the type of reaction

Since \(\Delta H = 560\space kJ>0\), the reaction is endothermic (heat is absorbed).

Step2: Calculate the moles of \(Fe_{2}O_{3}\)

The molar mass of \(Fe_{2}O_{3}\) is \(M=(2\times55.85 + 3\times16)=159.7\space g/mol\).
The number of moles \(n=\frac{m}{M}\), where \(m = 90.5\space g\). So \(n=\frac{90.5\space g}{159.7\space g/mol}\approx0.5667\space mol\)

Step3: Calculate the heat absorbed

From the reaction \(2Fe_{2}O_{3}(s)\to4FeO(s)+O_{2}(g)\space\Delta H = 560\space kJ\), for \(2\space mol\) of \(Fe_{2}O_{3}\), the heat change is \(560\space kJ\).
Let \(x\) be the heat absorbed for \(0.5667\space mol\) of \(Fe_{2}O_{3}\). Using the proportion \(\frac{x}{0.5667}=\frac{560}{2}\)
\(x=\frac{560\times0.5667}{2}=158.676\approx159\space kJ\)

Answer:

• This reaction is: endothermic
• Will any heat be released or absorbed: Yes, absorbed
• If heat is absorbed, the amount of heat: \(159\space kJ\)