Key Concepts of Chapter

1
(No Transcript)
2
Key Concepts of Chapter
Nature of Energy (units, kinetic and
potential). Transfer of Energy System and
Surroundings First Law of Thermodynamics
Internal Energy of a System in Relation to Heat
Endothermic vs. Exothermic Processes State
Functions Enthalpy and Heats of Reactions
Calorimetry Hesss Law Enthalpies of
Formation
3
Thermodynamics
the study of energy and its transformations
4
Thermochemistry
the study of the relationships between chemical
reactions and energy changes involving heat.
5
Energy
The capacity to do work or to transfer heat.
6
Kinetic Energy and Potential Energy
7
Kinetic energy
--energy of motion
What would an objects kinetic energy depend upon?
mass
Speed
8
Energy
Mass
Speed
9
What is the SI unit for mass?
What is the SI unit for distance?
What is the SI unit for time?
2
m
kg
x
________
2
s
10
Joule
2
m
kg
x
________
2
s
11
calorie
The amount of energy needed to raise the
temperature of 1 g of water 1 degrees Celsius.
1 cal 4.184 J
12
1 nutritional calorie
1 cal 4.184 J
Cal 1000 calories
13
Kinetic energy
14
Calculate the energy in joules of a 70.0 g
baseball moving at 40.0 m/s.
15
Convert this to energy in calories.
16
If we threw a water logged baseball at the same
speed, would the kinetic energy be higher or
lower?
17
What is the kinetic energy in joules of a 3.0
pound brick if it is tossed at 40.0 m/s?
1 Kg 2.2 lbs
18
Atoms and molecules have mass and are in motion
and therefore possess kinetic energy!
19
Potential energy
--stored energy
PE mgh
20
potential
kinetic
21
(No Transcript)
22
(No Transcript)
23
Forces other than gravity can cause potential
energy.
Electrostatic Force
24
Chemical Energy
due to the potential energy stored in the
arrangements of atoms
25
System
vs.
Surrounding
26
System
part of the universe we are interested in
studying.
Surrounding
the rest of the universe not involved in the
system
27
(No Transcript)
28
Bomb Calorimeter
Used to measure the heat released from combustion
reactions.
29
Coffee Cup Calorimter
Closed System Exchanges heat, but not matter with
surroundings.
30
Transferring Energy
Work
Heat
Energy used to move an object against a force
Energy transferred from a hotter object to a
colder object
F x d
31
W F x D
In physics, if you havent already, you will
learn about force.
F ma
Acceleration constant due to gravity Equal to 9.8
m/s2
How much work, in Joules, is needed to lift a box
weighing 10.0 Kg 3.15 meters?
32
m 10.0 Kg d 3.15 m a
9.8 m/s2
W F x d
W ma x d
W 314 J
33
Ethanol Combustion Reaction
Observe the following reaction. In your notes,
describe the reaction in terms of Potential
Energy Kinetic Energy System vs.
Surrounds Heat Work
34
First Law of Thermodynamics
In any chemical or physical change, energy cannot
be created nor destroyed, but it can be changed
in form.
(All energy lost by the system must be gained by
the surroundings, and vice versa.)
Energy System Energy
Surrounding
?Esys -?Esurr
35
Internal energy (E) is the sum of all energies
in the system (kinetic and potential)
Change in energy equals final energy minus
initial energy ?E Efinal - Einitial
Chemical Reaction Energy Einitial
reactants Efinal products
36
?E must include a number (magnitude), sign
(direction), and unit
Note We cannot directly determine initial and
final energies, but we can determine the change
in energy of a system.
37
system loses energy
2H2 (g) O2 (g) ? 2H2O
(l) energy
system gains
energy energy H2O (l ) ? 2H2 (g) O2 (g)
38
Internal Energy Diagram
H2(g), O2(g)
When hydrogen and oxygen react to form water,
energy is lost from the system.
?E lt 0
?E gt 0
Internal Energy, E
H2O(l)
39
?E, Heat Work

• Systems exchange energy with surroundings by heat
  work.
• When a system undergoes a physical or chemical
  change, the internal energy of the system
  changes.
• ?? q w
• q heat w work

40
We must consider the signs for q and w when doing
calculations involving the transfer of energy.
What happens to internal energy if the system
does work on the surroundings?
What happens to the internal energy if the system
loses heat?
decreases
decreases
?E q w Therefore we would expect q to be
negative if E decreases.
?E q w Therefore we would expect w to be
negative if E decreases.
If system loses heat q negative (-)
If system gains heat q positive ()
If system does work w negative (-)
If surroun does work w positive ()
41
Sample Exercise 5.3 from page 160 The hydrogen
and oxygen gases in the cylinder below are
ignited. As the reaction occurs, the system
loses 1150 J of heat to the surroundings. The
reaction also causes the piston to rise as the
hot gases expand. The expanding gas does 480 J
of work on the surroundings as it pushes up.
What is the change in the internal energy of the
system?
Heat released from system to surroundings q -
1150 J
?E q w
System does work on surroundings w - 480 J
?E q w ?E (- 1150 J) (- 480 J) -
1630 J
1630 J is transferred to the surroundings.
42
Exothermic Process
Process in which heat is released from the
system. Example Combustion reactions.
Endothermic Process
Process in which heat is absorbed by the
system. Example Ice melting.
43
Internal Energy is a state function.
State function property of a system that is
determined by it current state, and not how it
got to be in that state.
A very simple example Temperature of water at
room temperature (20-25C). The temperature could
have been achieved as a result of ice cold water
gaining heat or very hot water losing heat.
Either way, the state of the water at room
temperature is the same!
Properties that are state functions P, V, T, E,
H, S, G
PETS GO TO THE VET
Battery Handout
44
Enthalpy (H)

• Describes the heat content of a system.
• Enthalpy is a state function.
• Depends only on present state, not history.
• Cant measure H of a system.
• Can only measure ??.

45
Enthalpy of Reaction (Heat of reaction)
??rxn

• Because little work is done in chemical reactions
  (with the exception of gases), the ?H is related
  to heat.
• Heat of reactions describe the heat gained or
  lost under constant pressure.
• ?H Hfinal Hinitial qp
• ?H Hproducts - Hreactants

46
Enthalpy of Reaction (Heat of reaction)
??rxn

• Exothermic reactions have - ??
• 2H2(g) O2(g) ? 2H2O(g) ?? - 483.6kJ
• Endothermic reactions have ??
• 2H2O (l) ? 2H2O (g) ?? 88 KJ

47
(No Transcript)
48
(No Transcript)
49
Enthalpy Diagram
H 2(g) O2(g)
?H lt 0
Exothermic
Enthalpy
H2O (g)
50
Thermochemical Equations Balanced equations that
allow us to relate heat gained or lost by a
system to the of moles involved in a reaction.

N2 (g) 3H2 (g) ? 2NH3 ?H - 91.8 KJ Is this
endothermic or exothermic?
What if we doubled the amounts of reactants?
What would the change in enthalpy be? ?H -
183.6 KJ
What do you think the ?H would be if we reverse
the reaction? 2NH3 (g) ? N2 (g) 3H2 (g) ?H
91.8 KJ
51
Thermochemical equation guidelines

• Enthalpy is an extensive property (depends on
  amount).
• 2) Enthalpy change for a rx is equal in
  magnitude but opposite in sign to ?? for the
  reverse reaction.
• 3) Enthalpy change for a reaction depends on the
  state of the reactants and products.

52
Consider the following equations
CH4 (g) 2O2 (g) ? CO2 (g) 2H2O (l)
?H -890 KJ
CH4 (g) 2O2 (g) ? CO2 (g) 2H2O (g)
?H -802 KJ
Remember, the evaporation of water is
endothermic! 88 KJ of heat are needed to
evaporate 2 mol water
53
Time to calculate!
54
How many KJ of heat are released when 18.00 grams
of methane gas is burned in oxygen?
?H -890.0 KJ
CH4 (g) 2O2 (g) ? CO2 (g) 2H2O (l)
18.0 g CH4
1 mol CH4
-890.0 KJ
16.0 g CH4
1 mol CH4
-1001 KJ
55
Now you try one!
The complete combustion of 1 mol of butane gas,
is highly exothermic releasing 2845 KJ of heat.
Write the balanced thermochemical equation and
determine the enthalpy change if you burn 50.0
grams of butane.
2C4H10 (g) 13O2 (g) ? 8CO2 (g) 10H2O (g)
?H - 5690 KJ
50.0 g C4H10
1 mol C4H10
-5690 KJ
58.14 g C4H10
2 mol C4H10
-2447 KJ
56
Since we were given the enthalpy value of 1 mol
of butane, you could have left the balanced
equation in the form of 1 mol of butane.
C4H10 (g) 13/2O2 (g) ? 4CO2 (g) 5H2O (g)
?H - 2845 KJ
50.0 g C4H10
1 mol C4H10
-2845 KJ
58.14 g C4H10
1 mol C4H10
-2447 KJ
57
Calorimetry

• The measurement of heat flow (??) in a system.
• Done at a constant pressure
• Calorimeter is the instrument used for this
  measurement.

58
Heat Capacity
Amount of heat required to raise the temperature
of an object 1C.
The greater the heat capacity, the greater the
amount of heat needed to raise the temperature of
the object. ?? Which would have a higher heat
capacity, a sheet of aluminum or a sheet of black
poster board ?? Poster board because aluminum
would rise in temp quickly. Notice This term is
not related to any specific amount. It is simply
a term describing the general process.
59
When amounts are specified, it is usually in
terms of grams or moles.
Specific Heat Capacity
Amount of heat required to raise the temperature
of 1 gram of an object 1C.
Molar Heat Capacity
Amount of heat required to raise the temperature
of 1 mol of an object 1C.
Lets look at the math involved.
60

• q quantity of heat transferred
• m mass of substance in grams
• ?? temperature change (ºC or K)
• s often used as symbol for S.H.

How much heat is needed to raise 10.0g of
aluminum from 22ºC to 42ºC if the specific heat
of Al is 0.900 J/gK?
61
q 1.80 x 102 J
62
What is the molar heat capacity of aluminum? (We
are given the specific heat of aluminum in the
last problem, 0.900 J/gK. We must simply
convert the amount of grams into moles.)

• 24.3 J/molK

63
You try one!

• How much heat is needed to raise 0.0510 Kg of
  Iron from 23ºC to 33ºC if the specific heat of Fe
  is 0.450 J/gK?
• B) What is the molar heat capacity of Fe?

q 2.30 x 102 J
64
25.0 J/molK
65
Constant-Pressure Calorimetry

• The assumption is that all heat produced by a
  reaction is contained in the calorimeter.
• Exothermic aqueous rx
• Heat is lost by rx (or by metal) and gained by
  solution.
• Temperature of solution goes up, which can be
  measured.

66
Method we will use for Specific Heat Lab
67
Objective Determine the specific heat of Fe.
Heat 5.00 grams of Fe to 100C by placing it in
boiling water.Remove the Fe from the boiling
water and place it in a coffee-cup calorimeter
filled with 250.0 grams of room temperature water
(25.0C).Measure the change in temperature of
the water, which will equal the temperature
change of the Fe. The final temperature of the
water is 25.2C. These temperature changes are
related to the heat lost by the Fe and gained by
the water, so qm qH2O. The specific heat of
water is 4.184 J/gC.
Since qm qH2O, we can say that (Must
remember this step!!)
68
Given mm 5.00 grams ?Tm 100.0 C 25.2 C
74.8 C mH2O 250.0 grams SH2O 4.184 J/g
C ?TH2O 25.2 C 25.0 C .200 ºC ? Sm
69
Specific heat of Fe 0.559 J / g C
70
You try one!
The current room temperature is 23.0C. We place
a 3.00 gram piece of Al in a test tube into
boiling water (100.0 C). You then pour the test
tube of Al into a coffee cup calorimeter which
contains 100.0 grams of room temperature water
(25.0C). The temperature of the water rises and
stops rising at 23.497 C. What is the specific
heat of Al?
71
Specific heat of Al 0.906 J / g C
72
Consider the previous example. Instead of using
a 3.00 gram piece of aluminum, we use a 4.50 gram
piece of aluminum. Also, instead of using 100.0
grams of room temperature water, we use 50.0
grams of room temperature water. What
temperature will the water in the calorimeter
reach after placing the Al in the calorimeter?
Remember, we found that Specific heat of Al
0.906 J / g C.
73
?TH2O 1.50 C But we want to know the final
temp of the water.
23.0 C 1.50 C 24.5 C
Final Temp of water should 24.5 C
74

• Another Type of Problem
• When 50.0 mL of 0.10M AgNO3 is added to 50.0 mL
  of 0.10M HCl in a constant pressure calorimeter,
  the temperature increases from 25.0C to 25.8C.
• AgNO3(aq) HCl (aq) ? AgCl (s) HNO3 (aq)
• Calculate the ?? on a molar basis (Kj/mol) for
  this reaction assuming that the solution after
  the reaction has a density of 1.00 g/ml, and
  specific heat of 4.184 J/gK.

75
?H -qp
density mass/volume mass density X volume
s 4.184 J/g K ?T 0.800 C
76
334.7 J represents the amount of heat released.
But we are interested in the change in heat in Kj
per mol. We must consider the of moles of the
reactant that produced this amount of heat in
order to consider the enthalpy in Kj/mole.
mol M x L
mol 0.100 mol x 0.050 L 5.0 x 10-3 mol
L This can be
considered for either product, since the volumes
and molarity are the same for both. Also,
because the balanced equations shows a 11 ratio
for reactant/product, we can consider this of
moles for our product.
77
Now that we know how many moles were associated
with the of joules calculated, we can calculate
the enthalpy change in KJ/mol for the reaction.
1 KJ
-334.7 J
5.0 x 10-3 mol
1000 J
-66.9 KJ mol-1
78
Bomb Calorimetry (Constant volume)

• Combustion Reactions
• Organic compounds reacted with excess oxygen.
• The Heat Capacity of the calorimeter must be
  calibrated very precisely.

79
Hess's Law
If a reaction is carried out in a series of
steps, the ?H for the reaction is the sum of ?H
of each individual step.
Allows us to calculate ?H for many reactions
from a few experimentally determined
measurements. Allows us to calculate enthalpy
data for reactions which are difficult to carry
out directly.
80
Example Combustion of methane gas to form liquid
water may involve 2 steps.
CH4(g) 2O2(g) ? CO2(g) 2H2O(g) ?H -802 kj
2H2O(g) ? 2H2O(l) ?H -88 kj
__________________________________________________
_______________
CH4(g) 2O2(g) 2H2O(g) ? CO2(g) 2H2O(l)
2H2O(g)
?H -890 kj
Net equation CH4(g) 2O2(g) ? CO2(g) 2H2O(l)
81

• Because H is a state function, we will always get
  the same value for ?? no matter how many steps we
  take to get to the final product.

82
Using the following thermochemical equations,
calculate the ?H of combustion of C to CO.
(1) C(s) O2(g) ? CO2(g) ?H
-393.5 kj (2) CO(g) 1/2O2(g) ? CO2(g) ?H
-283.0 kj
(1) C(s) O2(g) ? CO2(g)
?H -393.5 kj
(2)CO(g) 1/2O2(g) ? CO2(g) ?H
-283.0 kj
(2) CO2(g) ? CO(g) 1/2O2(g) ?H 283.0
kj
__________________________________________________
C(s) 1/2O2(g)? CO(g) ?H -110.5 kj
83
Calculate ?H for the following reaction 2C(s)
H2(g) ?C2H2(g) given the following reactions
84
2C(s) H2(g) ?C2H2(g)
85
Enthalpies of Formation

• Heat involved with the process of elements
  chemically combining to form a compound.
• Ex C(s) O2(g) ? CO2(g) ?H -787.0 Kj/mol-1)
• ??f , where f indicates a substance formed from
  its elements.
• Reported in KJ/mol

86

• Magnitude of enthalpy depends on temperature,
  pressure, and state of reactants and products.
• To compare different rxs we need a standard.
• Standard enthalpy is defined as all reactants and
  products in their standard states. (??º)

87
Standard States T 298 K or 25C P 1
atmosphere If more than one state exists, the
most stable form is used. Examples
88
Standard Enthalpy of Formation

• ??ºf
• The change in enthalpy for the reaction that
  forms 1 mol of the compound from its elements,
  with all substances in standard state.

89
Many standard states which have been determined
experimentally are tabulated into data
tables. ?Hf Data Tables Page 177 (mini table
5.3), Appendix C of textbook, Values or Table
Given on AP Exam
90
(No Transcript)
91
May also see the following Enthalpies of
vaporization ?H for converting liquids to
gases Enthalpies of fusion ?H for melting
solids Enthalpies of combustion ?H for combusting
a substance with O2
Handout
92
hydrocarbons
(compounds containing only carbon and hydrogen)
Examples
ethane
C2H6
methane
propane
CH4
C3H8
93
Combustion of Hydrocarbons
Hydrocarbons react with oxygen to form carbon
dioxide and water.
C3H8(g) 5O2(g) ? 3CO2(g) 4H2O(l)
compounds containing
C
H
O
react with oxygen in the same manner
CH3OH(l) 3/2O2(g) ? CO2(g) 2H2O(l)
One of the few cases where you must leave
balanced equations in fraction form so we can
consider 1 mol values.
94
Write a balanced equation for the combustion of 1
mol of butane (C4H10).
C4H10(g) O2(g) ? CO2(g)
H2O(l)
Write a balanced equation for the combustion of 1
mol of ethane (C2H6).
C2H6(g) 7/2O2(g) ? 2CO2(g) 3H2O(l)
95
(sum of enthalpies of formation of
products) minus (sum of enthalpies of formation
of reactants)
Coefficients from equation
the sum of
96
Calculate the standard enthalpy change for the
combustion of 1 mol of benzene C6H6(l).
C6H6(l) 15/2O2(g) ? 6CO2(g) 3H2O(l)
- 3267 kj
97
Calculate the standard enthalpy change for the
combustion of 1 mol of ethanol(l).
C2H5OH(l) 3O2(g) ? 2CO2(g) 3H2O(l)
- 1367 kj