Step-by-step explanation:
Its value is 6.02 kj/mol. This means for every mole of ice we melt we must apply 6.02 kj of heat. We can calculate the heat needed with the following equation:
q
=
n
×
Δ
H
where:
q
= heat
n
= moles
Δ
H
= enthalpy
In this problem we would like to calculate the heat needed to melt 35 grams of ice at 0 °C. This problem can be broken into three steps:
1. Calculate moles of water
2. multiply by the enthalpy of fusion
3. Convert kJ to J
Step 1: Calculating moles of water
35
g
×
(
1
m
o
l
18.02
g
)
=
1.94
m
o
l
s
Step 2: Multiply by enthalpy of fusion
q
=
n
×
Δ
H
=
1.94
×
6.02
=
11.678
k
J
Step 3: Convert kJ to J
11.678
k
J
×
(
1000
J
1
k
J
)
=
11
,
678
J
Finally rounding to 2 sig figs (since 34°C has two sig figs) we get
q
=
12
,
000
J
One last note, if the temperature were not 0 °C then the ice would have to be heated in addition to melted. This would be a phase change problem combined with a heat capacity problem.