prentice hall © 2003chapter 5 chapter 5 thermochemistry chemistry the central science 9th edition...
Post on 21-Dec-2015
218 Views
Preview:
TRANSCRIPT
Prentice Hall © 2003 Chapter 5
Chapter 5Chapter 5ThermochemistryThermochemistry
CHEMISTRY The Central Science
9th Edition
David P. White
Prentice Hall © 2003 Chapter 5
Kinetic Energy and Potential Energy• Kinetic energy is the energy of motion:
• 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 EnergyThe Nature of Energy
2
21
mvEk
Prentice Hall © 2003 Chapter 5
Kinetic Energy and Potential Energy
• Electrostatic potential energy, Eel, is the attraction between two oppositely charged particles, Q1 and Q2, a distance d apart:
• The constant = 8.99 109 J-m/C2.
• If the two particles are of opposite charge, then Eel is the electrostatic attraction between them.
The Nature of EnergyThe Nature of Energy
Prentice Hall © 2003 Chapter 5
Units of Energy• SI Unit for energy is the joule, J:
We sometimes use the calorie instead of the joule:1 cal = 4.184 J (exactly)
A nutritional Calorie:1 Cal = 1000 cal = 1 kcal
The Nature of EnergyThe Nature of Energy
J 1m/s-kg 1
m/s 1kg 221
21
2
22
mvEk
Prentice Hall © 2003 Chapter 5
Systems and Surroundings• System: part of the universe we are interested in.• Surroundings: the rest of the universe.
The Nature of EnergyThe Nature of Energy
Prentice Hall © 2003 Chapter 5
Transferring Energy: Work and Heat• Force is a push or pull on an object.• Work is the product of force applied to an object over a
distance:
• Energy is the work done to move an object against a force.
• Heat (which doesn’t exist-Dr. Hansen) is the transfer of energy between two objects.
• Energy is the capacity to do work or transfer heat.
The Nature of EnergyThe Nature of Energy
dFw
Prentice Hall © 2003 Chapter 5
Internal Energy• Internal Energy: total energy of a system.• Cannot measure absolute internal energy.• Change in internal energy,
The First Law of The First Law of ThermodynamicsThermodynamics
initialfinal EEE
Prentice Hall © 2003 Chapter 5
Relating E to Heat and Work• 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:
when a system undergoes a physical or chemical change, the change in internal energy is given by the heat added to or absorbed by the system plus the work done on or by the system:
The First Law of The First Law of ThermodynamicsThermodynamics
wqE
Prentice Hall © 2003 Chapter 5
Exothermic and Endothermic Processes
• Endothermic: absorbs heat from the surroundings.• Exothermic: transfers heat to the surroundings.• An endothermic reaction feels cold.• An exothermic reaction feels hot.
The First Law of The First Law of ThermodynamicsThermodynamics
Prentice Hall © 2003 Chapter 5
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 The First Law of ThermodynamicsThermodynamics
State Functions
Prentice Hall © 2003 Chapter 5
• 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,
EnthalpyEnthalpy
VPw
EnthalpyEnthalpy
Prentice Hall © 2003 Chapter 5
• Enthalpy, H: Heat transferred between the system and surroundings carried out under constant pressure.
• Enthalpy is a state function.• If the process occurs at constant pressure,
EnthalpyEnthalpy
PVEH
VPE
PVEH
Prentice Hall © 2003 Chapter 5
• Since we know that
• We can write
• When H, is positive, the system gains heat from the surroundings.
• When H, is negative, the surroundings gain heat from the system.
EnthalpyEnthalpy
VPw
wq
VPEH
P
Prentice Hall © 2003 Chapter 5
EnthalpyEnthalpy
Prentice Hall © 2003 Chapter 5
• For a reaction:
• Enthalpy is an extensive property (magnitude H is
directly proportional to amount):
CH4(g) + 2O2(g) CO2(g) + 2H2O(l) H = -890 kJ
2CH4(g) + 4O2(g) 2CO2(g) + 4H2O(l) H = 1780 kJ
Enthalpies of ReactionEnthalpies of Reaction
reactantsproducts
initialfinal
HH
HHH
Prentice Hall © 2003 Chapter 5
• When we reverse a reaction, we change the sign of H:
CO2(g) + 2H2O(l) CH4(g) + 2O2(g) H = +890 kJ
• Change in enthalpy depends on state:
H2O(g) H2O(l) H = -88 kJ
Enthalpies of ReactionEnthalpies of Reaction
Prentice Hall © 2003 Chapter 5
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.
CalorimetryCalorimetry
Tq substance of gramsheat specific
Prentice Hall © 2003 Chapter 5
Constant Pressure Calorimetry• Atmospheric pressure is constant!
CalorimetryCalorimetry
T
qH P
solution of grams
solution ofheat specificsolnrxn
Prentice Hall © 2003 Chapter 5
CalorimetryCalorimetry
Constant Pressure Calorimetry
Prentice Hall © 2003 Chapter 5
CalorimetryCalorimetry
TCq calrxn
• Reaction carried out under constant volume.
• Use a bomb calorimeter.• Usually study combustion.
Bomb Calorimetry (Constant Volume Calorimetry)
Prentice Hall © 2003 Chapter 5
• 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
Hess’s LawHess’s Law
Prentice Hall © 2003 Chapter 5
Note that: H1 = H2 + H3
Hess’s LawHess’s Law
Prentice Hall © 2003 Chapter 5
Example 1.
Na(s) + 1/2Cl2(g) Na(g) + Cl(g) ; ΔH = +230kJ
Na(g) + Cl(g) Na+(g) + Cl-(g); ΔH = +147kJ
Na(s) + ½ Cl2(g) NaCl(s); ΔH = -411 kJ
Calculate ΔH for the reaction
Na+(g) + Cl-(g) NaCl(s)
Prentice Hall © 2003 Chapter 5
• If 1 mol of compound is formed from its constituent elements, then the enthalpy change for the reaction is called the enthalpy of formation, Ho
f .
• Standard conditions (standard state): 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 FormationEnthalpies of Formation
Prentice Hall © 2003 Chapter 5
• Example 2. Fe(s) + Br2(l) FeBr2(s); ΔH = -249.8 kJ
• FeBr3(s) FeBr2(s) + ½ Br2(l); ΔH = +18.4 kJ
• Calculate the heat of formation of FeBr3.
Prentice Hall © 2003 Chapter 5
• Answers • -147 -230 -411 kJ = -788 kJ• -249.8-18.4 kJ = -268.2 kJ
Prentice Hall © 2003 Chapter 5
• 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.
• Example – carbon’s most stable form is graphite.
oxygen’s most stable form is O2
Enthalpies of FormationEnthalpies of Formation
Prentice Hall © 2003 Chapter 5
Enthalpies of FormationEnthalpies of Formation
Prentice Hall © 2003 Chapter 5
Using Enthalpies of Formation of Calculate Enthalpies of Reaction
• We use Hess’ Law to calculate enthalpies of a reaction from enthalpies of formation.
Enthalpies of FormationEnthalpies of Formation
Prentice Hall © 2003 Chapter 5
Using Enthalpies of Formation of Calculate Enthalpies of Reaction
• For a reaction
Enthalpies of FormationEnthalpies of Formation
reactantsproductsrxn ff HmHnH
Prentice Hall © 2003 Chapter 5
Hrxn° = n Hf° (products) - n Hf° (rctnts)• n = moles• C. Ex: combustion of ethane • 2 C2H6 (g) + 7 O2 (g) --> 4 CO2 (g) +6 H2O(l) • H°rxn = n Hf° (prod) - n Hf° (rctnts) H°rxn = 4 mol Hf°CO2 + 6 mol Hf°H2O - 2 mol Hf° C2H6 -7 mol Hf°
O2
• = 4 mol (-393.51 kJ/mol) + 6 mol (-285.83 kJ/mol • - 2 mol (-84.68 kJ/mol) - 7 mol (O)• = -1574.0 kJ + (-1715.0 kJ) - (-169.4 kJ) - 0• = -3119.6 kJ• --> Note: 5 Steps to these problems
Prentice Hall © 2003 Chapter 5
Foods• Fuel value = energy released when 1 g of substance is
burned.• 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
Foods and FuelsFoods and Fuels
Prentice Hall © 2003 Chapter 5
Foods• Fats: contain more energy; are not water soluble, so are
good for energy storage.
Foods and FuelsFoods and Fuels
Prentice Hall © 2003 Chapter 5
Fuels• In 2002 the United States consumed 1.03 1017 kJ of
fuel.• Most from petroleum, natural gas and coal.• Remainder from nuclear, wind, and hydroelectric.• Fossil fuels are not renewable.
Foods and FuelsFoods and Fuels
Prentice Hall © 2003 Chapter 5
Foods and Foods and FuelsFuels
Prentice Hall © 2003 Chapter 5
Fuels• Fuel value = energy released when 1 g of substance is
burned.• Hydrogen has great potential as a fuel with a fuel value
of 142 kJ/g.
Foods and FuelsFoods and Fuels
top related