Transcript (Thermochemistry-Chapter 5) - Fall 2015

Thermochemistry
Energy
1st Law of Thermodynamics
Enthalpy / Calorimetry
Hess' Law
Enthalpy of Formation
The Nature of Energy
Kinetic Energy and Potential Energy
• Kinetic energy is the energy of motion:
1
Ek   m  v 2
2
• Potential energy is the energy an object possesses by
virtue of its position.
• Potential energy can be converted into kinetic energy.
Example: a bicyclist at the top of a hill.
The Nature of Energy
Units of Energy
• SI Unit for energy is the joule, J:
1
1
2
2
Ek   m  v   2 kg   1 m/s 
2
2
2 -2
 1 kg m s  1 J
sometimes the calorie is used instead of the joule:
1 cal = 4.184 J (exactly)
A nutritional Calorie:
1 Cal = 1000 cal = 1 kcal
Thermochemistry Terminology
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System: part of the universe we are interested in.
Surrounding: the rest of the universe.
Boundary: between system & surrounding.
Exothermic: energy released by system to surrounding.
Endothermic: energy absorbed by system from surr.
• Work ( w ): product of force applied to an
object over a distance.
•Heat ( q ): transfer of energy between two objects
The First Law of Thermodynamics
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Internal Energy
Internal Energy: total energy of a system.
Involves translational, rotational, vibrational motions.
Cannot measure absolute internal energy.
Change in internal energy,
E  Efinal  Einitial
DQ
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The First Law of Thermodynamics
Relating E to Heat(q) and Work(w)
• Energy cannot be created or destroyed.
• Energy of (system + surroundings) is constant.
• Any energy transferred from a system must be transferred
to the surroundings (and vice versa).
• From the first law of thermodynamics:
E  q  w
The First Law of
Thermodynamics
First Law of Thermodynamics
• Calculate the energy change for a system
undergoing an exothermic process in which
15.4 kJ of heat flows and where 6.3 kJ of
work is done on the system.
E = q + w
The First Law of Thermodynamics
Exothermic and Endothermic Processes
• Endothermic: absorbs heat from the surroundings.
• An endothermic reaction feels cold.
• Exothermic: transfers heat to the surroundings.
• An exothermic reaction feels hot.
Endothermic Reaction
Ba(OH)2•8H2O(s) + 2 NH4SCN(s) 
Ba(SCN)2(s) + 2 NH3(g) + 10 H2O(l)
The First Law of Thermodynamics
State Functions
• State function: depends only on the initial and final states
of system, not on how the internal energy is used.
The First Law of Thermodynamics
Enthalpy
• Chemical reactions can absorb or release heat.
• However, they also have the ability to do work.
• For example, when a gas is produced, then the gas
produced can be used to push a piston, thus doing work.
Zn(s) + 2H+(aq)  Zn2+(aq) + H2(g)
• The work performed by the above reaction is called
pressure-volume work.
• When the pressure is constant,
w   PV
Enthalpy
Enthalpy
• Enthalpy, H: Heat transferred between the system and
surroundings carried out under constant pressure.
H  E  PV
• Enthalpy is a state function.
• If the process occurs at constant pressure,
H   E  PV 
 E  PV
Enthalpy
• Since we know that
w   PV
• We can write
H  E  PV
 q P  w  PV
 q P  (  PV )  PV
 qP
• When H is positive, the system gains heat from the
surroundings.
• When H is negative, the surroundings gain heat from
the system.
Enthalpy => Heat of Reaction
Enthalpies of Reaction
• For a reaction:
H  H final  H initial
 H products  H reactants
• Enthalpy is an extensive property (magnitude H is
directly proportional to amount):
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
H = -802 kJ
2CH4(g) + 4O2(g)  2CO2(g) + 4H2O(g) H = 1604 kJ
Enthalpies of Reaction
• When we reverse a reaction, we change the sign of H:
CO2(g) + 2H2O(g)  CH4(g) + 2O2(g)
H = +802 kJ
• Change in enthalpy depends on state:
H2O(g)  H2O(l) H = -44 kJ
Calorimetry
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Heat Capacity and Specific Heat
Calorimetry = measurement of heat flow.
Calorimeter = apparatus that measures heat flow.
Heat capacity = the amount of energy required to raise
the temperature of an object (by one degree).
Molar heat capacity = heat capacity of 1 mol of a
substance.
Specific heat = specific heat capacity = heat capacity of 1
g of a substance.
q  specific heat  grams of substance  T
Table 5.2: Specific Heats (S) of Some Substances at 298 K
Substance
N2(g)
Al(s)
Fe(s)
Hg(l)
H2O(l)
H2O(s)
CH4(g)
CO2(g)
Wood , Glass
S ( J g-1 K-1 )
1.04
0.902
0.45
0.14
4.184
2.06
2.20
0.84
1.76 , 0.84
If 24.2 kJ is used to warm a piece of aluminum with a mass of 250. g, what is
the final temperature of the aluminum if its initial temperature is 5.0oC?
q  S m T
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Calorimetry
Constant Pressure Calorimetry
• Atmospheric pressure is constant!
H  qP
qrxn  qsoln  specific heat of solution 
 grams of solution   T
Calorimetry
Constant Pressure Calorimetry
qrxn  S  ( total mass of solution)  T
H rxn
qrxn

mol *
* moles of species of interest
Aqueous Solution Assumptions:
Calorimetry
Calorimetry Examples
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1. In an experiment similar to the procedure set out for Part
(A) of the Calorimetry experiment, 1.500 g of Mg(s) was
combined with 125.0 mL of 1.0 M HCl. The initial
temperature was 25.0oC and the final temperature was
72.3oC. Calculate: (a) the heat involved in the reaction
and (b) the enthalpy of reaction in terms of the number of
moles of Mg(s) used. Ans: (a) –25.0 kJ (b) –406 kJ/mol
2. 50.0 mL of 1.0 M HCl at 25.0oC were mixed with 50.0
mL of 1.0 M NaOH also at 25.0oC in a styrofoam cup
calorimeter. After the mixing process, the thermometer
reading was at 31.9oC. Calculate the energy involved in
the reaction and the enthalpy per moles of hydrogen ions
used. Ans: -2.9 kJ , -58 kJ/mol [heat of neutralization
for strong acid/base reactions]
In an experiment similar to the procedure set out for Part (A) of the Calorimetry experiment, 1.500 g of
Mg(s) was combined with 125.0 mL of 1.0 M HCl. The initial temperature was 25.0oC and the final
temperature was 72.3oC. Calculate: (a) the heat involved in the reaction and (b) the enthalpy of reaction
in terms of the number of moles of Mg(s) used.
qrxn  S  ( total mass of solution)  T
qrxn= - S x msoln x ∆T
∆H = -406 kJ/mol
2. 50.0 mL of 1.0 M HCl at 25.0oC were mixed with 50.0 mL of 1.0 M NaOH also at 25.0oC in a styrofoam cup
calorimeter. After the mixing process, the thermometer reading was at 31.9 oC. Calculate the energy involved in the
reaction and the enthalpy per moles of hydrogen ions used. Ans: -2.9 kJ , -58 kJ/mol [heat of neutralization for
strong acid/base reactions]
Hess’s Law
• Hess’s law: if a reaction is carried out in a number of
steps, H for the overall reaction is the sum of H for
each individual step.
• For example:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
H = -802 kJ
2H2O(g)  2H2O(l)
H= - 88 kJ
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)
H = -890. kJ
Another Example of Hess’s Law
Given:
C(s) + ½ O2(g)  CO(g)
H = -110.5 kJ
CO2(g)  CO(g) + ½ O2(g)
H = 283.0 kJ
Calculate H for:
C(s) + O2(g)  CO2(g)
Enthalpies of Formation
• If 1 mol of compound is formed from its constituent
elements, then the enthalpy change for the reaction is
called the enthalpy of formation, Hof .
• Standard conditions (standard state): Most stable form of
the substance at 1 atm and 25 oC (298 K).
• Standard enthalpy, Ho, is the enthalpy measured when
everything is in its standard state.
• Standard enthalpy of formation: 1 mol of compound is
formed from substances in their standard states.
Enthalpies of Formation
• If there is more than one state for a substance under
standard conditions, the more stable one is used.
• Standard enthalpy of formation of the most stable form of
an element is zero.
Enthalpies of Formation
Substance
Hof (kJ/mol)
C(s, graphite)
0
O(g)
247.5
O2(g)
0
N2(g)
0
Bucky Ball drawn by HyperChem
Enthalpies of Formation
Using Enthalpies of Formation to
Calculate Enthalpies of Reaction
• For a reaction
H rxn   n  H f products    m  H f reactants 
• Note: n & m are stoichiometric coefficients.
• Calculate heat of reaction for the combustion of propane
gas giving carbon dioxide and water.
C3H8(g) + 5O2(g)  3CO2(g) + 4H2O()
C3H8(g) + 5O2(g)  3CO2(g) + 4H2O()
∆Hof (kJ/mol):
C3H8(g) + 5O2(g)  3CO2(g) + 4H2O()
∆Hof (kJ/mol):
∆Horxn = -2,219.9 kJ/mol
Foods and Fuels
Foods
• 1 nutritional Calorie, 1 Cal = 1000 cal = 1 kcal.
• Energy in our bodies comes from carbohydrates and fats
(mostly).
• Intestines: carbohydrates converted into glucose:
C6H12O6 + 6O2  6CO2 + 6H2O, H = -2816 kJ
• Fats break down as follows:
2C57H110O6 + 163O2  114CO2 + 110H2O, H = -75,520 kJ
Fats contain more energy; are not water soluble, so are good for
energy storage.
Foods and Fuels
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Fuels
Fuel value = energy released when 1 g of substance is
burned.
Most from petroleum and natural gas.
Remainder from coal, nuclear, and hydroelectric.
Fossil fuels are not renewable.
In 2000 the United States consumed 1.03  1017 kJ of
fuel.
• In 2005 the United States consumed 1.05  1017 kJ of energy.
• In 2008 the United States consumed 1.05x1017 kJ of energy.
• Hydrogen has great potential as a fuel with a fuel value of
142 kJ/g. [ gasoline ≈ 35 kJ/g ]
[8.5%]
[7.4%]
Foods and
Fuels
[37.5%]
[22.6%]
(% for 2000)
[24.0%]
[% for 2008]
Energy in Fuels
The vast
majority of the
energy
consumed in this
country comes
from fossil fuels.
© 2012 Pearson Education, Inc.
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Thermochemistry
Ek 
1
 m  v2
2
E  q  w
(kg m 2 s  2  joule)
Energy
1st Law of Thermodynamics
H  qP
qrxn  qsoln  specific heat of solution  grams of solution  T
Enthalpy / Calorimetry
Hess' Law
Enthalpy of Formation
H rxn   n  H f products    m  H f reactants 