2. Energy Capacity to do work or to produce heat.
Law of conservation of energy – energy can be converted from one form to another but can be neither created nor destroyed.
The total energy content of the universe is constant.
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3. Energy Potential energy – energy due to position or composition.
Kinetic energy – energy due to motion of the object and depends on the mass of the object and its velocity.
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4. Energy
Heat involves the transfer of energy between two objects due to a temperature difference.
Work – force acting over a distance.
Energy is a state function; work and heat are not
State Function – property that does not depend in any way on the system’s past or future (only depends on present state).
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5. Chemical Energy
System – part of the universe on which we wish to focus attention.
Surroundings – include everything else in the universe.
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6. open
closed
isolated
Exchange:
mass & energy
energy
nothing
6.2

7. Chemical Energy Endothermic Reaction: Heat flow is into a system.
Absorb energy from the surroundings.
Exothermic Reaction:
Energy flows out of the system.
Energy gained by the surroundings must be equal to the energy lost by the system.
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8. CONCEPT CHECK!
Is the freezing of water an endothermic or exothermic process? Explain.
Exothermic process because you must remove energy in order to slow the molecules down to form a solid.
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9. CONCEPT CHECK!
Classify each process as exothermic or endothermic. Explain. The system is underlined in each example.
Your hand gets cold when you touch ice.
The ice gets warmer when you touch it.
Water boils in a kettle being heated on a stove.
Water vapor condenses on a cold pipe.
Ice cream melts.
Exo
Endo
Exothermic (heat energy leaves your hand and moves to the ice)
Endothermic (heat energy flows into the ice)
Endothermic (heat energy flows into the water to boil it)
Exothermic (heat energy leaves to condense the water from a gas to a liquid)
Endothermic (heat energy flows into the ice cream to melt it)
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10. Methane is burning in a Bunsen burner in a laboratory.
CONCEPT CHECK!
For each of the following, define a system and its surroundings and give the direction of energy transfer.
Methane is burning in a Bunsen burner in a laboratory.
Water drops, sitting on your skin after swimming, evaporate.
System – methane and oxygen to produce carbon dioxide and water; Surroundings – everything else around it; Direction of energy transfer – energy transfers from the system to the surroundings (exothermic)
System – water drops; Surroundings – everything else around it; Direction of energy transfer – energy transfers from the surroundings (your body) to the system (water drops) (endothermic)
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11. Thermodynamics
The study of energy and its interconversions is called thermodynamics.
Law of conservation of energy is often called the first law of thermodynamics.
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12. Internal Energy
Internal energy E of a system is the sum of the kinetic and potential energies of all the “particles” in the system.
To change the internal energy of a system: ΔE = q + w
q represents heat
w represents work
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13. Internal Energy Thermodynamic quantities consist of two parts:
Number gives the magnitude of the change.
Sign indicates the direction of the flow.
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14. Internal Energy Sign reflects the system’s point of view.
Endothermic Process:
q is positive
Exothermic Process:
q is negative
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15. Internal Energy Sign reflects the system’s point of view.
System does work on surroundings:
w is negative
Surroundings do work on the system:
w is positive
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16. Work
For an expanding gas, ΔV is a positive quantity because the volume is increasing. Thus ΔV and w must have opposite signs:
w = –PΔV
To convert between L·atm and Joules, use 1 L·atm = J.
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17. They perform the same amount of work.
EXERCISE!
Which of the following performs more work?
a) A gas expanding against a pressure of 2 atm from 1.0 L to 4.0 L.
A gas expanding against a pressure of 3 atm from 1.0 L to 3.0 L.
They perform the same amount of work.
They both perform the same amount of work. w = -PΔV
a) w = -(2 atm)( ) = -6 L·atm
b) w = -(3 atm)( ) = -6 L·atm
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18. Calculate the change in internal energy of the system when it undergoes an exothermic process releasing 140 J of heat and a gas is evolved which does 85 J of work during expansion.
A balloon is heated by adding 240 J of heat. It expands doing 135 J of work. Calculate the change in internal energy of the balloon.
50.0 g of Fe(s) is cooled from 100C to 90C losing 225 J of heat. Calculate the change in internal energy of the Fe(s)
They both perform the same amount of work. w = -PΔV
a) w = -(2 atm)( ) = -6 L·atm
b) w = -(3 atm)( ) = -6 L·atm
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19. Change in Enthalpy State function ΔH = q at constant pressure
ΔH = Hproducts – Hreactants
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20. –252 kJ Consider the combustion of propane: EXERCISE! ΔH = –2221 kJ
C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(l)
ΔH = –2221 kJ
Assume that all of the heat comes from the combustion of propane. Calculate ΔH in which 5.00 g of propane is burned in excess oxygen at constant pressure.
–252 kJ
(5.00 g C3H8)(1 mol / g C3H8)(-2221 kJ / mol C3H8)
ΔH = -252 kJ
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21. Calorimetry Science of measuring heat Specific heat capacity:
The energy required to raise the temperature of one gram of a substance by one degree Celsius.
Molar heat capacity:
The energy required to raise the temperature of one mole of substance by one degree Celsius.
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22. Calorimetry Heat change = s × m × ΔT
s = specific heat capacity (J/°C·g)
m = mass of system in units of g
ΔT = change in temperature (°C)
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23. The molar heat capacity of Al is 24. 4J/mol-oC
The molar heat capacity of Al is 24.4J/mol-oC. Calculate its specific heat capacity. Answer: J/g-oC
The specific heat capacity of Hg is 0.14J/g-oC. Calculate the amount of heat required to raise the temperature of 10.0g of Hg from 25.0C to 30.0C. Answer: 7.0J
The specific heat of copper is J/g°C. Calculate its molar heat capacity. Answer: 24.5 J/mol-oC
If J of heat is supplied to 10.0g of Copper, what is the temperature change of the metal? The specific heat of copper is J/g°C. Answer: 26.0 °C

24. A Coffee–Cup Calorimeter Made of Two Styrofoam Cups
qrxn = - qsolution =-m x s x Dt
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25. When 50. 0 mL of 0. 100M AgNO3(aq) and 50. 0 mL of 0
When 50.0 mL of 0.100M AgNO3(aq) and 50.0 mL of M HCl(aq) are mixed in a coffee cup calorimeter, the temperature increased from 22.30C to 23.11C.
AgNO3(aq) + HCl(aq)  HNO3(aq) + AgCl(s)
Calculate the enthalpy change for the reaction per mol HCl.
The density of the solution is 1.03 g/mL. Its specific heat = 4.18J/gC.

26. When 4. 25 g of NH4NO3 is dissolved in 60
When 4.25 g of NH4NO3 is dissolved in 60.0g of water in a coffe-cup calorimeter, the temperature drops from 22.0C to 16.9C. Calculate the enthalpy change, H in kJ/mol NH4NO3 . The specific heat of the solution = J/gC.

27. Constant-Volume Calorimetry qrxn = - qcalorimeter
qcalorimeter = Ccalorimeter x Dt
Reaction at Constant V
DH = qrxn= DE
DH ~ qrxn
No heat enters or leaves!
6.5

28. When 4. 25 g of NH4NO3 is dissolved in 60
When 4.25 g of NH4NO3 is dissolved in 60.0g of water in a coffe-cup calorimeter, the temperature drops from 22.0C to 16.9C. Calculate the enthalpy change, H in kJ/mol NH4NO3 . The specific heat of the solution = J/gC.

29. When a hot object and a cold object are brought in contact,
qcold = -qhot
qcold = heat change of the cold object
qhot = heat change of the hot object
Heat flows from the hot to the cold till they are both at the same final temperature, equilibrium temperature, T
qcold = mcold x scold x (T-Ti,cold)
qhot = mhot x shot x (T-Ti,hot)

30. CONCEPT CHECK!
A g sample of water at 90°C is added to a g sample of water at 10°C. The final temperature of the water is: a) Between 50°C and 90°C b) 50°C c) Between 10°C and 50°C
The correct answer is b).
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31. CONCEPT CHECK!
A g sample of water at 90.°C is added to a g sample of water at 10.°C. The final temperature of the water is: a) Between 50°C and 90°C b) 50°C c) Between 10°C and 50°C Calculate the final temperature of the water. 23°C
The correct answer is c). The final temperature of the water is 23°C.
- (100.0 g)(4.18 J/°C·g)(Tf – 90.) = (500.0 g)(4.18 J/°C·g)(Tf – 10.)
Tf = 23°C
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32. CONCEPT CHECK!
You have a Styrofoam cup with 50.0 g of water at 10.°C. You add a 50.0 g iron ball at 90. °C to the water. (sH2O = 4.18 J/°C·g and sFe = 0.45 J/°C·g) The final temperature of the water is: a) Between 50°C and 90°C b) 50°C c) Between 10°C and 50°C Calculate the final temperature of the water. 18°C
The correct answer is c). The final temperature of the water is 18°C.
- (50.0 g)(0.45 J/°C·g)(Tf – 90.) = (50.0 g)(4.18 J/°C·g)(Tf – 10.)
Tf = 18°C
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33. In going from a particular set of reactants to a particular set of products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps.
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34. N2(g) + 2O2(g) → 2NO2(g) ΔH1 = 68 kJ
This reaction also can be carried out in two distinct steps, with enthalpy changes designated by ΔH2 and ΔH3.
N2(g) + O2(g) → 2NO(g) ΔH2 = 180 kJ
2NO(g) + O2(g) → 2NO2(g) ΔH3 = – 112 kJ
N2(g) + 2O2(g) → 2NO2(g) ΔH2 + ΔH3 = 68 kJ
ΔH1 = ΔH2 + ΔH3 = 68 kJ
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35. The Principle of Hess’s Law
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36. Characteristics of Enthalpy Changes
If a reaction is reversed, the sign of ΔH is also reversed.
The magnitude of ΔH is directly proportional to the quantities of reactants and products in a reaction. If the coefficients in a balanced reaction are multiplied by an integer, the value of ΔH is multiplied by the same integer.
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37. Example Consider the following data: Calculate ΔH for the reaction
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38. Problem-Solving Strategy
Work backward from the required reaction, using the reactants and products to decide how to manipulate the other given reactions at your disposal.
Reverse any reactions as needed to give the required reactants and products.
Multiply reactions to give the correct numbers of reactants and products.
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39. Example Reverse the two reactions: Desired reaction:
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40. Example
Multiply reactions to give the correct numbers of reactants and products:
4( ) 4( )
3( ) 3( )
Desired reaction:

41. Example Final reactions: Desired reaction: ΔH = +1268 kJ
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42. Standard Enthalpy of Formation (ΔHf°)
Change in enthalpy that accompanies the formation of one mole of a substance from its elements with all substances in their standard states.
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43. Conventional Definitions of Standard States
For a Compound
For a gas, pressure is exactly 1 atm.
For a solution, concentration is exactly 1 M.
Pure substance (liquid or solid)
For an Element
The form [N2(g), K(s)] in which it exists at 1 atm and 25°C.
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44. Problem-Solving Strategy: Enthalpy Calculations
The change in enthalpy for a given reaction can be calculated from the enthalpies of formation of the reactants and products:
ΔH°rxn = ΣnpΔHf°(products) - ΣnrΔHf°(reactants)
Elements in their standard states are not included in the ΔHreaction calculations because ΔHf° for an element in its standard state is zero.
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45. 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g) Given the following information:
EXERCISE!
Calculate ΔH° for the following reaction:
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
Given the following information:
ΔHf° (kJ/mol)
Na(s) 0
H2O(l) –286
NaOH(aq) –470
H2(g) 0
ΔH° = –368 kJ
[2(–470) + 0] – [0 + 2(–286)] = –368 kJ
ΔH = –368 kJ
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