Chapter 12, Problem 95E

Interpretation Introduction

Interpretation: Calculate the amount of heat evolved when 1.00 mol of steam at $145°C$ is converted to ice at $-50.0°C$.

Concept Introduction: Heat of fusion is the amount of heat required to melt 1 mole of a solid. It is denoted by the symbol $\text{Δ}{H}_{\text{fus}}$. $\text{Δ}{H}_{\text{fus}}$ for water is equal to $6.02\text{ kJ/mol}$.

Heat of vaporization is the amount of heat required to vaporize 1 mole of liquid.

It is denoted by the symbol $\text{Δ}{H}_{\text{vap}}$.

$\text{Δ}{H}_{\text{vap}}$ for water is equal to $\text{40}\text{.7 kJ/mol}$.

Heat released by water $\left(q\right)$ while changing its temperature is calculated as:

$q=mC\Delta T$

Here, m is the mass of water, C is the heat capacity of water in that state and $\Delta T$ is the temperature change.

Heat capacity of liquid water is $4.186\text{ J}/\text{g }°\text{C}$.

Heat capacity of steam is $\text{1}\text{.84 J}/\text{g }°\text{C}$.

Heat capacity of ice is $\text{2}\text{.09 J}/\text{g }°\text{C}$.

The expression of number of moles $\left(n\right)$ in terms of molar mass is:

$n=\frac{\text{mass}}{\text{molar mass}}$