# Question #3de10

Jul 7, 2017

$\text{29,900 J}$

#### Explanation:

In order to be able to answer this question, you need to know the value of the enthalpy of fusion, $\Delta {H}_{\text{fus}}$, of water, which you'll find listed as

$\Delta {H}_{\text{fus" = "333.55 J g}}^{- 1}$

https://en.wikipedia.org/wiki/Enthalpy_of_fusion

Now, the enthalpy of fusion of a given substance tells you the amount of heat needed in order to convert $\text{1 g}$ of said substance from solid at its melting point to liquid at its melting point.

In this case, you can say that in order to convert $\text{1 g}$ of ice at its normal melting point of ${0}^{\circ} \text{C}$ to $\text{1 g}$ of liquid water at ${0}^{\circ} \text{C}$, you need to provide it with $\text{333.55 J}$ of heat.

You can thus say that you will need

$89.5 \textcolor{red}{\cancel{\textcolor{b l a c k}{\text{g"))) * overbrace("333.55 J"/(1color(red)(cancel(color(black)("g")))))^(color(blue)(=DeltaH_"fus")) = color(darkgreen)(ul(color(black)("29,900 J}}}}$

in order to melt $\text{89.5 g}$ of ice at its normal melting point.

The answer is rounded to three sig figs, the number of significant figures you have for the mass of ice.