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Calorimetry: Heat of reaction
Related Topics
Wize High School Grade 12 Chemistry Textbook > Energy Changes
Calorimetry
6 Activities
The temperature of a 12.58 g sample of calcium carbonate (CaCO
3
(s)) increases from 23.6°C to 38.2°C. If the specific heat capacity of calcium carbonate is 0.82 J/g K, how many joules of heat are absorbed?
151 J
5.0 J
7.5 J
410 J
0.82 J
I don't know
Check Submission
More Calorimetry Questions:
Heat of Vaporization for Hexane
Hexane is heated from 0
o
C to 50
o
C and the equilibrium vapour pressure increases from 0.06 atm to 0.53 atm.
Calculate the Δ
vap
H for hexanes.
Thermochemistry: Melting ice
The standard enthalpy of fusion for water,
Δ
H
fusion
o
=
6.01
kJ/mol
.
\Delta H^o_{\textrm{fusion}}=6.01\ \textrm{kJ/mol}.
Δ
H
fusion
o
=
6.01
kJ/mol
.
Calculate the energy required (in KJ) to convert 26.0 g of ice at 0
o
C to liquid water at 20
o
C.
Calorimetry: Final temperature calculation
The ΔH
soln
for the process when solid sodium hydroxide (NaOH) dissolves in water is 44 kJ/mol. When a 10.0 g sample of NaOH dissolves in 250.0 g of water in a coffee-cup calorimeter, the temperature increases from 23.0
o
C to _______
o
C. Assume that the solution has the same specific heat as liquid water, i.e. 4.18 J/gK.
Calorimetry: Heat of combustion
When 2g of rocket fuel, N
2
H
4
is burned inside of a bomb-calorimeter, the temperature of the water rises from 22
o
C to 29
o
C. If there is 0.9kg of water and the heat capacity for the bomb is 822J/
o
C, what is the heat of the combustion for one mole of N
2
H
4
in the bomb calorimeter?
Thermochemistry: Enthalpy change
The enthalpy of vaporization for water, ΔH
o
vaporization
= 40.7 kJ/mol. Determine the enthalpy change (in kJ) when 54.0 mL of water is condensed from steam at 100
o
C.
Thermochemistry: Melting ice
The standard enthalpy of fusion for water,
Δ
H
fusion
o
=
6.01
kJ/mol
.
\Delta H^o_{\textrm{fusion}}=6.01\ \textrm{kJ/mol}.
Δ
H
fusion
o
=
6.01
kJ/mol
.
Calculate the energy required (in KJ) to convert 26.0 g of ice at 0
o
C to liquid water at 20
o
C.
Calorimetry: Heat of combustion
When 2g of rocket fuel, N
2
H
4
is burned inside of a bomb-calorimeter, the temperature of the water rises from 22
o
C to 29
o
C. If there is 0.9kg of water and the heat capacity for the bomb is 822J/
o
C, what is the heat of the combustion for one mole of N
2
H
4
in the bomb calorimeter?
Calorimetry: Heat of combustion
When 2g of rocket fuel, N
2
H
4
is burned inside of a bomb-calorimeter, the temperature of the water rises from 22
o
C to 29
o
C. If there is 0.9kg of water and the heat capacity for the bomb is 822J/
o
C, what is the heat of the combustion for one mole of N
2
H
4
in the bomb calorimeter?
Enthalpy of Formation
What is the enthalpy of the following reaction when 0.35 g of Mg
2+
is produced? Use the following information:
ΔH°
f
[Mg(OH)
2
(s)] = −924.5 kJ
ΔH°
f
[NH
4+
(aq)] = −132.5 kJ
Thermodynamics: Internal energy and Enthalpy
2.3 moles of an ideal gas initially at -12
o
C and 16 atm underwent a complete expansion against a constant external pressure of 1atm and did 1.1kJ of work. Calculate Δ𝑈, Δ𝐻, 𝑎𝑛𝑑 𝑞 for this process.
Calorimetry: Final temperature calculation
The Δ
H
soln
for the process when solid sodium hydroxide (NaOH) dissolves in water is 44 kJ/mol. When a 10.0 g sample of NaOH dissolves in 250.0 g of water in a coffee-cup calorimeter, the temperature increases from 23.0
o
C to _______
o
C. Assume that the solution has the same specific heat as liquid water ie, 4.18 J/gK.
Calorimetry: Heat of reaction
0.23 g of ethanol (C
2
H
5
OH) was burned in a bomb calorimeter (constant
volume
) leading to a temperature change from 22.50
o
C to 23.81
o
C. If the heat capacity of the calorimeter is 5.2 kJ/K, what would be the amount of heat generated by the combustion of exactly one mole of ethanol? ___________
Calorimetry: Temperature change
1. A 56.0 g block of iron heated to 90 °C is placed in a 4.8 kg water bath at 20 °C. What will be the final temperature of the iron? heat capacity of iron c = 0.450 J /gK, heat capacity of water is 4.185 J/gK
Calorimetry: Temperature change
How much heat energy is required to raise the temperature of 35.0 g of Al (c = 0.902 J/g K) from 18 °C to 65 °C? If this much energy were instead used to raise the temperature of 2.0 L of water, by how much would the temperature of the water change? Round to the nearest
o
C.
Calorimetry: Heat of reaction
4. A bomb calorimeter assembly has a heat capacity of 650 J/°C. The assembly is placed in a 3.0 L water bath and 2.5g of a fuel compound is burned in excess oxygen. If the temperature of the water bath increases by 4.8 °C, what is the value of q
rxn
?
What is the amount of heat released per gram of compound burned?
Calorimetry: Heat capacity
5. A fuel compound is burned in a bomb calorimeter, releasing 458 kJ of heat. The temperature of the calorimeter assembly and surrounding 2.5 L water bath increases from 23.5 °C to 28.6 °C. What is the heat capacity of the calorimeter assembly?
Calorimetry
When 1.00 kg of lead (specific heat =0.13 J g
-1
°C
-1
) at 125
o
C is added to a quantity of water at 18.0°C, the final temperature of the lead-water mixture is 32.0°C. What is the volume of water present?
Note: the density of H
2
O
(l)
is 1.00g mL
-1
and the specific heat capacity is 4.18J g
-1 o
C
-1
A 2.00 mole sample of a monoatomic ideal gas is initially at 1.00 L and 1.00 atm. Calculate the total q, w,
Δ
E
{\small \Delta\text{E}}
Δ
E
and
Δ
H
{\small \Delta \text{H}}
Δ
H
when the final conditions are 0.500 L and 5.00 atm along the pathway where the volume is compressed first.
[Do NOT include units in your answers]
Thermochemistry: Energy required to melt ice
You are lost on an expedition and you need some water to drink. There is ice around at an ambient temperature of -22
o
C. Using the thermodynamic information below, how much energy is required to take 1 kg of ice and warm it up to water at 12
o
C.
c
p
,
s
(
i
c
e
)
=
2.01
J
K
−
1
m
o
l
−
1
Δ
H
f
u
s
o
=
6.01
k
J
m
o
l
−
1
c
p
,
s
(
w
a
t
e
r
)
=
4.18
J
K
−
1
m
o
l
−
1
Δ
H
f
u
s
o
=
44.00
k
J
m
o
l
−
1
\begin{array}{lr} c_{p,s}(ice) = 2.01\ J K^{-1} mol^{-1} & \Delta H^o_{fus} = 6.01\ kJ mol^{-1} \\ c_{p,s}(water) = 4.18\ J K^{-1} mol^{-1} & \Delta H^o_{fus} = 44.00\ kJ mol^{-1} \end{array}
c
p
,
s
(
i
ce
)
=
2.01
J
K
−
1
m
o
l
−
1
c
p
,
s
(
w
a
t
er
)
=
4.18
J
K
−
1
m
o
l
−
1
Δ
H
f
u
s
o
=
6.01
k
J
m
o
l
−
1
Δ
H
f
u
s
o
=
44.00
k
J
m
o
l
−
1
Heat of Vaporization for Hexane
Hexane is heated from 0
o
C to 50
o
C and the equilibrium vapour pressure increases from 0.06 atm to 0.53 atm.
a) Calculate the Δ
vap
H for hexanes.
b) The human nose can detect hexane at a partial pressure of 0.005 atm. Estimate the temperature at which we can no longer smell hexane.
Copper, Silver, water ,Gold, hydrogensulfide and have specific heats of 0.385, 0.235, 4.18, 0.129, and 1.342 Jg
-1o
C
-1
respectively. Which of the following would need smallest quantity of heat to change the temperature of 5400kg by 65
o
C
An alkane’s vapour pressure at 25.00 °C is 6.00 kPa. What is its vapour pressure at 80.00 °C if the enthalpy of vaporization is 43.0 kJ/mol? (Give your answer in kPa)
Sucrose, C
12
H
22
O
11
, can be metabolized for energy in the body according to the reaction: C
12
H
22
O
11
(s) + 12 O
2
(g)
→
\rightarrow
→
12 CO
2
(g) + 11 H
2
O (l) ΔH
o
= -5.65x10
3
kJ
How many grams of sucrose must be burned to raise the temperature of 50.0 kg of water by 1.00
o
C? The specific heat capacity of water is 4.184 J g
-1
K
-1
.
Calorimetry: Energy of vaporization of water
How much energy is required to bring 1.00 g of liquid water from 35 °C to 165 °C water vapour? The enthalpy of vaporization of water is +40.7 kJ/mol, the specific heat capacity of liquid water is 4.18 J / g °C, and the specific heat capacity of water vapour is 1.87 J/g °C.
From the following reaction:
2
C
6
H
12
O
6
(
s
)
+
2
O
2
(
g
)
→
12
C
O
2
(
g
)
+
4
H
2
O
(
l
)
Δ
H
r
x
n
=
−
1680
k
J
2C_6H_{12} O_{6(s)}+2O_{2(g)} \to 12CO_{2(g)}+4H_2O_{(l)} \hspace{20pt} \Delta H_{rxn}=-1680 kJ
2
C
6
H
12
O
6
(
s
)
+
2
O
2
(
g
)
→
12
C
O
2
(
g
)
+
4
H
2
O
(
l
)
Δ
H
r
x
n
=
−
1680
k
J
When an unknown quantity of sugar (C
6
H
12
O
6
) is burned, it was determined that the internal energy of the system decreased by 960 kJ and that 15.7 kJ of work was done on the system. How many grams of sugar were present at the beginning of this specific reaction?
A chemical compound has a molecular weight of 89.05 g/mol. 1.400 grams of this compound underwent complete combustion under constant pressure conditions in a calorimeter with a heat capacity of 2.980 × 10
3
J/ °C. The temperature went up by 11.95 °C. Calculate the standard heat of combustion of the compound.
Calorimetry: Heat capacity of calorimeter
The combustion of 1.00 mol of glucose (C
6
H
12
O
6
) gives off 2.82x10
3
kJ of heat (i.e. ΔH
comb
= 2.82x10
3
kJ mol-1). If 1.80 g glucose (molar mass is 180) is burned in a bomb calorimeter and the temperature of the calorimeter increases from 19.00
o
C to 23.50
o
C, what is the heat capacity of the calorimeter (kJ/
o
C)?
Calorimetry: Heat requirement
How much heat is required to convert a 15.5 g ice cube at -5.00
o
C to liquid water at 0
o
C? (The specific heat of ice is 2.09 J/g K, the specific heat of water is 4.18 J/g K, and the ΔH
fus
for water=334 J/g).
Thermochemistry: Melting ice
The standard enthalpy of fusion for water,
Δ
H
fusion
o
=
6.01
kJ/mol
.
\Delta H^o_{\textrm{fusion}}=6.01\ \textrm{kJ/mol}.
Δ
H
fusion
o
=
6.01
kJ/mol
.
Calculate the energy required (in KJ) to convert 26.0 g of ice at 0
o
C to liquid water at 20
o
C.
Enthalpy: Heat of vaporization
The enthalpy of vaporization for water, ΔH
o
vaporization
= 40.7 kJ/mol. Determine the enthalpy change (in kJ) when 54.0 mL of water is condensed from steam at 100
o
C.
Calorimetry: Using Heat Qualitatively
Consider the following four samples, all initially at room temperature:
The glass cubes are cooled to the same low temperature, and each is placed into one of the containers of room temperature gas.
(a) If Container A had a lower final temperature at thermal equilibrium than Container B, which gas has the lower specific heat capacity?
Calorimetry
From the following reaction:
2
C
6
H
12
O
6
(
s
)
+
2
O
2
(
g
)
→
12
C
O
2
(
g
)
+
4
H
2
O
(
l
)
Δ
H
r
x
n
=
−
1680
k
J
/
m
o
l
⋅
r
x
n
2C_6H_{12}O_6(s)+2O_2(g)→12CO_2(g)+4H_2O(l)\qquad \qquad ΔH_{rxn}=-1680kJ/mol\cdot rxn
2
C
6
H
12
O
6
(
s
)
+
2
O
2
(
g
)
→
12
C
O
2
(
g
)
+
4
H
2
O
(
l
)
Δ
H
r
x
n
=
−
1680
k
J
/
m
o
l
⋅
r
x
n
When an unknown quantity of sugar (C
6
H
12
O
6
) is burned, it was determined that the internal energy of the system decreased by 960 kJ and that 15.7 kJ of work was done on the system. How many grams of sugar were present at the beginning of this specific reaction?
Thermochemistry: Calorimetry
Consider the following balanced chemical equation:
2
C
ℓ
F
3
(
g
)
+
2
O
2
(
g
)
→
C
ℓ
2
O
(
g
)
+
3
F
2
O
(
g
)
∆
r
H
°
=
?
?
?
2 CℓF_3(g) + 2 O_2(g) → Cℓ_2O (g) + 3 F_2O (g) \qquad ∆_rH° = ???
2
C
ℓ
F
3
(
g
)
+
2
O
2
(
g
)
→
C
ℓ
2
O
(
g
)
+
3
F
2
O
(
g
)
∆
r
H
°
=
???
S
u
b
s
t
a
n
c
e
△
f
H
°
@
25
°
C
(
k
J
/
m
o
l
)
C
ℓ
F
3
(
g
)
?
?
C
ℓ
2
O
(
g
)
80.3
F
2
O
(
g
)
–
21.7
F
2
(
g
)
0
\def\arraystretch{1.5} \begin{array}{|c|c|} \hline \rm Substance & △_fH° \ @ \ 25°C \ (kJ/mol) \\ \hline CℓF_3 (g) & ?? \\ \hline Cℓ_2O (g) & 80.3 \\ \hline F_2O (g) & – 21.7 \\ \hline F_2 (g) & 0 \\ \hline \end{array}
Substance
C
ℓ
F
3
(
g
)
C
ℓ
2
O
(
g
)
F
2
O
(
g
)
F
2
(
g
)
△
f
H
°
@
25°
C
(
k
J
/
m
o
l
)
??
80.3
–21.7
0
Thermochemistry: Calorimetry
Ethene gas (C
2
H
4
) reacts with fluorine gas according to the following balanced equation:
C
2
H
4
(
g
)
+
6
F
2
(
g
)
→
2
C
F
4
(
g
)
+
4
H
F
(
g
)
Δ
r
H
°
=
−
3012
k
J
/
m
o
l
⋅
r
x
n
@
25
°
C
C_2H_4 (g) + 6 F_2 (g) → 2 CF_4 (g) + 4 HF (g) \qquadΔ_rH° = −3012\ kJ/mol·rxn\ @ \ 25°C
C
2
H
4
(
g
)
+
6
F
2
(
g
)
→
2
C
F
4
(
g
)
+
4
H
F
(
g
)
Δ
r
H
°
=
−
3012
k
J
/
m
o
l
⋅
r
x
n
@
25°
C
S
u
b
s
t
a
n
c
e
△
f
H
°
@
25
°
C
(
k
J
/
m
o
l
)
C
2
H
4
(
g
)
?
?
F
2
(
g
)
0
C
F
4
(
g
)
–
933
H
F
(
g
)
–
273.3
\def\arraystretch{1.5} \begin{array}{|c|c|} \hline \rm{Substance} & △_fH° \ @ \ 25°C(kJ/mol) \\ \hline C_2H_4 (g) & ?? \\ \hline F_2 (g) & 0 \\ \hline CF_4 (g) & – 933 \\ \hline HF (g) & – 273.3 \\ \hline \end{array}
Substance
C
2
H
4
(
g
)
F
2
(
g
)
C
F
4
(
g
)
H
F
(
g
)
△
f
H
°
@
25°
C
(
k
J
/
m
o
l
)
??
0
–933
–273.3
Calculate ΔH
o
f
of SO
2(g)
given the following: (answer in kJ)
S
(
s
)
+
1.5
O
2
(
g
)
→
S
O
3
(
g
)
,
Δ
H
=
−
395.8
k
J
S_{(s)}+1.5O_{2(g)}\to SO_{3(g)}, \Delta H=-395.8kJ
S
(
s
)
+
1.5
O
2
(
g
)
→
S
O
3
(
g
)
,
Δ
H
=
−
395.8
k
J
2
S
O
2
(
g
)
+
O
2
(
g
)
→
2
S
O
3
(
g
)
,
Δ
H
=
−
198.2
k
J
2SO_{2(g)}+O_{2(g)}\to 2SO_{3(g)}, \Delta H=-198.2kJ
2
S
O
2
(
g
)
+
O
2
(
g
)
→
2
S
O
3
(
g
)
,
Δ
H
=
−
198.2
k
J
The standard enthalpy of formation for the amino acid glycine C
2
H
5
NO
2
(s) is –528 kJ/mol at 25°C. Which reaction below has an enthalpy change equal to this value?
Calorimetry
Equal masses of copper and iron are both heated to a temperature of 100°C. Each piece of hot metal is then placed into its own beaker containing 1 L of water. If both beakers start with the same initial temperature, which beaker of water will have the higher final temperature?
Additional information: Cu (specific heat = 0.377 J/g°C); Fe (specific heat = 0.450 J/g°C)
Calorimetry
The temperature of a 12.58 g sample of calcium carbonate (CaCO
3
(s)) increases from 23.6
o
C to 38.2
o
C. If the specific heat of calcium carbonate is 0.82 J/g K, how many joules of heat are absorbed?
Thermochemistry Calculations
The standard enthalpy of fusion for water, ΔH
o
fusion
= 6.01 kJ/mol. Calculate the energy required (in kJ) to convert 26.0 g of ice at 0
o
C to liquid water at 20
o
C.
Thermochemistry: Enthalpy change
The enthalpy of vaporization for water, ΔH
o
vaporization
= 40.7 kJ/mol. Determine the enthalpy change (in kJ) when 54.0 mL of water is condensed from steam at 100
o
C.
Bomb Calorimetry
A fuel compound was burned in oxygen in a bomb calorimeter with a combined heat capacity of 56.5 k] / K (Bomb assembly and water). The temperature of the surrounding water bath increased from 23.5 to 34.6 °C as a result. Calculate
Δ
E
s
y
s
\Delta E_{sys}
Δ
E
sy
s
for this process? Consider any change in pressure due to gaseous products to be negligible.