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How much energy is required to melt these quantities of ice?a. 56 $\mathrm{g}$b. 56 $\mathrm{mol}$

a.4474.4 $\mathrm{cal}$b. 80539 $\mathrm{cal}$

Chemistry 102

Chemistry 101

Chapter 18

Observing Energy

Section 6

Heat and Phase Changes

Thermodynamics

Thermochemistry

Carleton College

University of Maryland - University College

Brown University

Lectures

03:07

A liquid is a nearly incom…

04:38

A liquid is a state of mat…

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How much energy does it ta…

05:01

What mass of ice can be me…

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How much energy in kilojou…

08:33

What quantity of energy do…

07:13

01:49

The energy required to mel…

00:24

How much heat would it tak…

03:31

Calculate the quantity of …

02:09

The heat energy required t…

01:46

question Number six is a heat of fusion calculation. We recognize this because it asks us to determine how much energy is required to melt these quantities of ice and melting has a heat of fusion associated with it. Fusion is melting, so if we have 56 g of water and we recognize that the heat of fusion of, well, ice, we have 56 g of ice and the heat of fusion of ice water is 79.9 calories per gram, then all we need to do is multiply these two numbers together to get 4474 calories. But if we have 56 moles of ice, first we have to convert the moles of ice into grams of ice by multiplying by the molar mass of water, which is 18.1 g per mole. Then we can convert the grams into calories, using the heat of fusion of ice, 79.9 calories per gram, and we get 80,600 calories

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