Transcript q = nC∆T
Thermodymanics
Thermodynamics
is a branch of science that
focuses on energy changes that accompany
chemical and physical changes.
Objective:
To calculate heat capacity.
Heat
(q): The energy transferred between
objects that are at different temperatures.
Unit:
joules (J)
Molar
Heat Capacity: Energy (heat) needed
to increase the temperature of 1 mol of
substance by 1 K.
q = nC∆T
q = heat
n = # of moles
C = molar heat capacity
∆T = change in temperature
The
molar heat capacity of water is larger
than the molar heat capacity of land. This
means that water does not heat up as easily
as land does. As a result, oceans can help
keep coastal areas cool during the summer.
The filling of a fruit pie has a larger heat
capacity than the crust. This means that
fruit filling will retain heat better and the
crust will cool much quicker. As a result,
eating the fruit filling can cause burns (even
though it may appear that the pie is cool).
Determine
the energy (heat) needed to
increase the temperature of 10.0 mol of Hg
by 7.5 K. The value of C for Hg is 27.8
J/K۰mol.
Determine
the energy (heat) needed to
increase the temperature of 10.0 mol of Hg
by 7.5 K. The value of C for Hg is 27.8
J/K۰mol.
q=?
n = 10.0 mol
C = 27.8 J/K۰mol
∆T = 7.5 K
q = nC∆T
q = (10.0 mol)(27.8 J/K۰mol)(7.5 K)
q= 2.1 x 103 J
The
molar heat capacity of tungsten is 24.2
J/K۰ mol. Calculate the energy as heat
needed to increase the temperature of 0.404
mol of W by 10.0 K.
The
molar heat capacity of tungsten is 24.2
J/K۰ mol. Calculate the energy as heat
needed to increase the temperature of 0.404
mol of W by 10.0 K.
q=?
n = 0.404 mol
C = 24.2 J/K۰mol
∆T = 10.0 K
q = nC∆T
q = (0.404 mol)(24.2 J/K۰mol)(10.0 K)
q= 97.8J
Suppose
a sample of NaCl increased in temperature
by 2.5 K when the sample absorbed 1.7 x 102 J
energy (heat). Calculate the number of moles of
NaCl if the molar heat capacity is 50.5 J/K۰mol.
Suppose
a sample of NaCl increased in temperature
by 2.5 K when the sample absorbed 1.7 x 102 J
energy (heat). Calculate the number of moles of
NaCl if the molar heat capacity is 50.5 J/K۰mol.
q = 1.7 x 102 J
n=?
C = 50.5 J/K۰mol
∆T = 2.5 K
q = nC∆T
1.7 x 102 J = n(50.5 J/K۰mol)(2.5 K)
n= 1.3 mol
Calculate
the energy as heat needed to increase
the temperature of 0.80 mol of nitrogen, N2, by 9.5
K. The molar heat capacity of nitrogen is 29.1
J/K۰mol.
Calculate
the energy as heat needed to increase
the temperature of 0.80 mol of nitrogen, N2, by 9.5
K. The molar heat capacity of nitrogen is 29.1
J/K۰mol.
q=?
n = 0.80 mol
C = 29.1 J/K۰mol
∆T = 9.5 K
q = nC∆T
q = (0.80 mol)(29.1 J/K۰mol)(9.5 K)
q= 2.2 x 102 J
A
0.07 mol sample of octane, C8H18, absorbed 3.5 x
103 J of energy. Calculate the temperature
increase of octane if the molar heat capacity of
octane is 254.0 J/K۰mol.
A
0.07 mol sample of octane, C8H18, absorbed 3.5 x
103 J of energy. Calculate the temperature
increase of octane if the molar heat capacity of
octane is 254.0 J/K۰mol.
q = 3.5 x 103 J
n = 0.07 mol
C = 254.0 J/K۰mol
∆T = ?
q = nC∆T
3.5 x 103 J = (0.07 mol)(254.0 J/K۰mol) ∆T
∆T = 2.0 x 102 K
Objective:
enthalpy.
To calculate the change in
The
total energy of a system is impossible to
measure.
However, we can measure the change in
enthalpy of a system.
Enthalpy: The total energy content of a
sample.
H
= Enthalpy
∆H = Change in Enthalpy
∆H
A
can be measured with a calorimeter.
calorimeter is used to measure the heat
absorbed or released in a chemical or
physical change.
Enthalpy
changes can be used to determine if a
process is endothermic or exothermic.
Exothermic
Reaction: Negative Enthalpy
Change
Endothermic
Changes
Reaction: Positive Enthalpy
The
standard enthalpy of formation ( H f ) is
the enthalpy change in forming 1 mol of a
substance from elements in their standard
states.
0
Note:
The value for the standard enthalpy of
formation for an element is 0.
The
values for the standard enthalpies of
formation can be found using a table.
ΔHreaction = ΔH products - ΔHreactants
Step
1: Determine
for each compound
using enthalpy table.
Step 2: Multiply by the coefficients from the
balanced equation (# of moles).
Step 3: Set up ΔH equation.
Step 4: calculate.
H
0
f
ΔHreaction = ΔH products - ΔHreactants
Calculate ΔH for the following reaction
and determine if the reaction is
exothermic or endothermic.
SO2(g) + NO2(g) SO3(g) + NO(g)
ΔHreaction = ΔH products - ΔHreactants
SO2(g) + NO2(g) SO3(g) + NO(g)
-296.8(1)
33.1(1)
kJ/mol
kJ/mol
-395.8(1)
kJ/mol
90.3(1)
kJ/mol
ΔH= [(-395.8 kJ/mol)(1) + (90.3 kJ/mol)(1)] –
[(296.8 kJ/mol)(1) + (33.1 kJ/mol)(1)]
Δ H = -305.5 kJ – 329.9 kJ
Δ H = -635.4 kJ The reaction is exothermic.
Calculate
the enthalpy change for the
following reaction:
2C2H6(g) + 7O2(g) 4 CO2(g) + 6H2O(g)
Calculate
ΔH for the following reaction:
CaO(s) + H2O(l) Ca(OH)2(s)
Calculate
the enthalpy change for the
combustion of methane gas.
CH4(g) + 2O2(g) CO2(g) + 2H2O(l)
Objective:
entropy.
To calculate the change in
A
reaction is more likely to occur if enthalpy
(ΔH) is negative.
However, some endothermic reactions can
occur easily. Why? Entropy!
Entropy (ΔS): A measure of the randomness
or disorder of a system.
A process if more likely to occur if there is an
increase in entropy (or if ΔS is positive).
Entropy
is increased by the following factors:
Diffusion (process of dispersion)
Dilution of a solution
Decreasing the pressure of a gas
Increasing temperature
The number of moles of product is greater than
the number of moles of reactant
Increasing the total number of particles in a
system
When a reaction produces more gas particles
(opposed to liquid or solids)
ΔSreaction = ΔSproducts - ΔSreactants
Find
the change in entropy for the following
reaction:
2Na (s) + 2HCl (g) 2NaCl (s) + H2(g)
Find
the change in entropy for the following
reaction:
2Na (s) + 2H2O (l) 2NaOH (s) + H2(g)
Objectives:
(1)
(2)
To calculate the change in Gibbs energy.
To determine if a reaction is spontaneous
or nonspontaneous.
Gibbs
Energy: the energy in a system that is
available to do useful work.
Gibbs
ΔG
Energy is also called Free Energy.
= Change in Gibbs Energy
ΔGreaction = ΔGproducts - ΔGreactants
Calculate
ΔG for the following water-gas
reaction:
C(s) + H2O (g) CO (g) + H2 (g)
Calculate
the Gibbs energy change that
accompanies the following reaction:
C(s) + O2 (g) CO2 (g)
Calculate
the Gibbs energy change that
accompanies the following reaction:
CaCO3 (s) CaO (s) + CO2 (g)
If
ΔG is negative, the forward reaction is
spontaneous.
If
ΔG is 0, the system is at equilibrium.
If
ΔG is positive, the forward reaction is
nonspontaneous.
NOTE: In this case, the reaction is spontaneous in the
reverse direction.
ΔG = ΔH - TΔS
Step
1: Organize the information.
Step 2: Change the units.
Step 3: Step up ΔG equation.
Step 4: Calculate.
Given
the changes in enthalpy and entropy are -139
kJ and 277 J/K respectively for a reaction at 25⁰C,
calculate the change in Gibbs energy.
Given
the changes in enthalpy and entropy are -139
kJ and 277 J/K respectively for a reaction at 25⁰C,
calculate the change in Gibbs energy.
ΔH = -139 kJ
ΔS = 277 J/K / (1000 J/kJ) = 0.277 kJ/K
T = 25 ⁰C + 273 = 298 K
ΔG = ?
Given
the changes in enthalpy and entropy are -139
kJ and 277 J/K respectively for a reaction at 25⁰C,
calculate the change in Gibbs energy.
ΔH = -139 kJ
ΔS = 277 J/K / (1000 J/kJ) = 0.277 kJ/K
T = 25 ⁰C + 273 = 298 K
ΔG = ?
ΔG = ΔH - T ΔS
ΔG = (-139 kJ) – (298K)(0.277 kJ/K)
ΔG = (-139 kJ) – (82.546 kJ)
ΔG = -221.55 kJ The reaction is spontaneous.
A
reaction has a ΔH of -76 kJ and a ΔS of -117
J/K. Calculate the ΔG at 298 K. Is the reaction
spontaneous?
A
reaction has a ΔH of 11 kJ and a ΔS of 49 J/K.
Calculate ΔG at 298 K. Is the reaction
spontaneous?
Objective:
To calculate the melting and
boiling point of a substance.
T mp
Tmp:
H
fus
S fus
T bp
H vap
S vap
melting point temperature
Tbp: boiling point temperature
ΔHfus: molar enthalpy of fusion
ΔSfus: molar entropy of fusion
ΔHvap: molar enthalpy of vaporization
ΔSvap: molar entropy of vaporization
Step
1: Organize the information.
Step 2: Change the units.
Step 3: Step up the equation.
Step 4: Calculate.
The enthalpy of fusion of Hg is 2.295 kJ/mol, and the molar
entropy of fusion is 9.79 J/mol·K. The enthalpy of vaporization at
the boiling point if 59.2 kJ/mol, and the molar entropy of
vaporization is 93.8 J/mol·K. Calculate the melting point and boiling
point of Hg.
The enthalpy of fusion of Hg is 2.295 kJ/mol, and the molar
entropy of fusion is 9.79 J/mol·K. The enthalpy of vaporization at
the boiling point if 59.2 kJ/mol, and the molar entropy of
vaporization is 93.8 J/mol·K. Calculate the melting point and boiling
point of Hg.
Tmp = ?
ΔHfus = 2.295 kJ/mol
ΔSfus = 9.79 J/mol·K = 0.00979 kJ/mol·K
The enthalpy of fusion of Hg is 2.295 kJ/mol, and the molar
entropy of fusion is 9.79 J/mol·K. The enthalpy of vaporization at
the boiling point if 59.2 kJ/mol, and the molar entropy of
vaporization is 93.8 J/mol·K. Calculate the melting point and boiling
point of Hg.
Tmp = ?
ΔHfus = 2.295 kJ/mol
ΔSfus = 9.79 J/mol·K = 0.00979 kJ/mol·K
Tmp = (2.295 kJ/mol) / (0.00979 kJ/mol·K)
Tmp = 234 K
The enthalpy of fusion of Hg is 2.295 kJ/mol, and the molar
entropy of fusion is 9.79 J/mol·K. The enthalpy of vaporization at
the boiling point if 59.2 kJ/mol, and the molar entropy of
vaporization is 93.8 J/mol·K. Calculate the melting point and boiling
point of Hg.
Tmp = ?
ΔHfus = 2.295 kJ/mol
ΔSfus = 9.79 J/mol·K = 0.00979 kJ/mol·K
Tmp = (2.295 kJ/mol) / (0.00979 kJ/mol·K)
Tmp = 234 K
Tbp = ?
ΔHvap = 59.2 kJ/mol
ΔSvap = 93.8 J/mol·K = 0.0938 kJ/mol·K
The enthalpy of fusion of Hg is 2.295 kJ/mol, and the molar
entropy of fusion is 9.79 J/mol·K. The enthalpy of vaporization at
the boiling point if 59.2 kJ/mol, and the molar entropy of
vaporization is 93.8 J/mol·K. Calculate the melting point and boiling
point of Hg.
Tmp = ?
ΔHfus = 2.295 kJ/mol
ΔSfus = 9.79 J/mol·K = 0.00979 kJ/mol·K
Tmp = (2.295 kJ/mol) / (0.00979 kJ/mol·K)
Tmp = 234 K
Tbp = ?
ΔHvap = 59.2 kJ/mol
ΔSvap = 93.8 J/mol·K = 0.0938 kJ/mol·K
Tbp = (59.2 kJ/mol) / (0.0938 kJ/mol·K )
T = 631 K
For
ethanol, the molar enthalpy of fusion is
4.931 kJ/mol and the molar entropy of fusion
is 31.6 J/mol·K. The molar enthalpy of
vaporization at the boiling point is 42.32
kJ/mol and the molar entropy of
vaporization is 109.9 J/mol·K. Calculate the
melting and boiling points for ethanol.
For
sulfur dioxide, the molar enthalpy of
fusion is 8.62 kJ/mol and the molar entropy
of fusion is 43.1 J/mol·K. The enthalpy of
vaporization at the boiling point is 24.9
kJ/mol and the molar entropy of
vaporization is 94.5 J/mol·K. Calculate the
melting and boiling points for sulfur dioxide.
For
ammonia, ΔHfus is 5.66 kJ/mol and Δsfus is
29.0 J/mol·K. ΔHvap is 23.33 kJ/mol and ΔSvap is
97.2 J/mol·K. Calculate the melting and
boiling points for ammonia.