Chapter 5 Chemistry Lecture
Transcript of Chapter 5 Chemistry Lecture
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Energy
Types of energy
– Kinetic
– Potential
• Gravitational
• Electrostatic
• Chemical Measured in joules (J), or calories (cal)
1 cal = 4.184 J
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Systems and Surroundings
The system is what we are studying
The surroundings are everything else
Universe = system + surroundings
Types of systems
– Closed - common in chemistry• Exchanges energy but not mass with
surroundings
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What Energy Can Do
Energy can be gained by a system if:
– Work is done on the system
– Heat is gained by the system Energy can be lost by a system if:
– The system does work
– The system loses heat So, what is work?
For that matter, what is heat?
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Who’s Doing The Work?
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Question
If a student compresses a gas by pushingdown on a piston:
a. The student has done work on the system;energy of the system decreases
b. The system has done work on the student;energy of the system decreases
c. The student has done work on the system;
energy of the system increasesd. The system has done work on the student;
energy of the system increases
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First Law of Thermodynamics
Energy can neither be created nor
destroyed
An easier way to say this:
– Energy is conserved
An even easier way to say this:
– YOU CAN’T WIN!
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What Does This Mean?
Internal energy is the energy of a system (E)
If a system changes its energy, the new
energy is: – ∆E = Efinal - Einitial
∆E is negative if system has lost energy…to? – The surroundings!
∆E is positive if system has gainedenergy…from? – The surroundings!
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Put Another Way…
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Chemical Reactions and
Energy
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E, Heat, and Work
Systems and surroundings may exchangeenergy as – Heat
– Work
∆E = q + w – q = heat
• +q is heat GAINED by system, -q is heat LOST bysystem
– w = work• +w is work done ON system, -w is work done BY system
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Example
What is the ∆E for a system for a
process in which the system absorbs
140 J of heat from the surroundings,and does 85 J of work on the
surroundings?
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Heat
Endothermic (heat in) processes
– The system absorbs heat
Exothermic (heat out) processes
– The system loses heat
Temperature is proportional to the
internal energy of a system
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Exothermic and Endothermic
Processes
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State Functions
The total energy (E) for a system is hard
to know
– We generally talk about change in energy
Factors influencing E
– Temperature
– Pressure
– Total quantity of matter
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State Functions
Properties of systems that are defined
by the system’s condition (state) are
state functions – State functions depend upon conditions,
NOT pathway
– Internal energy is a state function
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Not State Functions
Heat (q) is not a state function – A system can get from one state to another
through loss/gain of various amounts of heat
Work (w) is not a state function – A system can get from one state to another
through doing/being subjected to various amountsof work
For a given energy change, q and w can vary,but E does not
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Reactions and Work
Most reactions in chemistry are open to
the atmosphere
– Pressure is constant
Work is associated with a change in
volume
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Question
Which of the following represents the
relationship between system and
surroundings energy?a. ∆Esystem = ∆Esurroundings
b. ∆Esystem = -∆Esurroundings
c. ∆Esystem = 2(∆Esurroundings)d. ∆Esystem = -2(∆Esurroundings)
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Question
Which of the following is a state
function?
a. Internal energy (of a system)
b. Heat
c. Work
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PV Work
PV work
– Change in gas volume
– W = -P∆V • Expansion: ∆V is positive, work BY system
• Compression: ∆V is negative, work ON system
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PV Work
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An Important Note
Temperature is proportional to the internal
energy of the system
– All other factors held constant, temperatureincreases if heat (energy) is absorbed
– All other factors held constant, temperature
increases if work is done ON system
– All other factors held constant, temperaturedecreases if work is done BY system
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Question
For gas in a piston canister, which of
the following would NOT constitute an
increase in E of the system?a. Pushing down on the piston
b. Lighting a candle under the canister
c. Pulling up on the piston
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Question
What does it mean to say temperature of asystem is proportional to E?
a. The temperature of the system can be used tocalculate E
b. Comparing the temperatures of differentsystems allows you to compare their Es
c. If ∆E for a system is positive, the temperature
will decreased. If ∆E for a system is positive, the temperature
will increase
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Enthalpy (H)
Provides a quantification of heat flow if
– Constant pressure
– No work other than PV work
– H = E + PV
• Internal energy plus product of
pressure/volume
• H is a state function
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∆H
∆H is change in enthalpy
∆H = ∆(E + PV), or ∆E + P∆V
But…easier…
– In systems at constant pressure, ∆H = q
• If ∆H is positive, ENDOTHERMIC reaction
• If ∆H is negative, EXOTHERMIC reaction
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Enthalpies of Reaction
For a reaction, ∆Hrxn = ∆Hprod - ∆Hreact
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Other Enthalpy Issues
What happens if we reverse the
reaction?
– Magnitude of ∆Hrxn is the same, signchanges
What happens if we burn two moles of
methane? – Twice as much heat is evolved
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Definition
Squir-rel-y adj. Slang
– 1.Eccentric.
– 2.Cunningly unforthcoming or reticent.
– The American Heritage Dictionary of the
English Language, Fourth Edition.
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Question
A squirrel is sitting on an electrical
wire. What kind of energy does he
have?a. Gravitational potential
b. Electrical potential
c. Kinetic
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Question
The squirrel leaps for a neighboring
wire; as he moves through the air
(quite some distance above theground), what kind of energy does he
have?
a. Gravitational potentialb. Kinetic
c. Gravitational potential and kinetic
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Calorimetry
How can we measure ∆Hrxn experimentally?
Need to know a few facts… – As substances gain heat, they get hotter (and vice
versa)
– Substances increase/decrease in temperature
predictably• Can quantify heat gained/lost based upon temperaturechange
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Heat Capacity (C)
The amount of heat necessary to raisetemperature of an object by 1 ºC (1 K)
– Larger heat capacity means more heatrequired to raise temperature
Molar heat capacity (Cmolar ) is heatcapacity of one mole of a substance
Specific heat (s) is heat capacity of onegram of a substance
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Specific Heat (s)
s = q
m x ∆T
Quantity of heat
transferred
Change in
temperatureGrams of
substanceCan also write in terms of heat:
q = ms∆T
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Example
What is the specific heat of water, given
that it takes 418 J of heat to raise the
temperature of 50.0 g by 2.00 K?
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Question
By what amount will the temperature of
50.0 g of water change if it absorbs 815
J of heat, given that swater = 4.18 J/gK?
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The “Coffee Cup” Calorimeter
Not sealed• Reaction occurs at
constant (atm) pressure
• Coffee cup is insulating,has low heat capacity, so
low heat absorption
• qsoln = (ssoln)(gsoln)(∆T) = -qrxn
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Example
50.0 mL 0.100 M AgNO3 and 50.0 mL
0.100 M HCl are mixed in a coffee cup
calorimeter. The temperature of thesolution increases from 22.20ºC to
23.11ºC; calculate enthalpy of reaction
for the reaction: AgNO3 (aq) + HCl (aq) AgCl (s) + HNO3 (aq)
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“Bomb” Calorimetry
Combustion in
sealed “bomb”
Heat absorbed bycontents of
calorimeter
qrxn = -Ccal∆T
Constant volume
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Example
A 1.800 g sample of phenol (MW =
94.11 g/mol) was burned in a bomb
calorimeter with a heat capacity of11.66 kJ/ºC. The temperature of the
calorimeter and contents increased from
21.36ºC to 26.37ºC. What is the heat ofcombustion per mole of phenol?
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Question
If a system changes temperature a
LOT for the amount of heat it absorbs,
what kind of specific heat does it have:a. High
b. Low
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Question
In a coffee-cup calorimeter, what do
we assume comprises the system to
SIMPLIFY our calculations?a. Everything in the universe
b. The cup plus the water in the cup
c. The water in the cup ONLY
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Hess’s Law
Very formal definition
– If a reaction is carried out in a series of
steps, ∆H for the overall reaction will equalthe sum of the ∆Hs for individual steps
What this really means
– We can use known ∆H values to determineunknown ∆H values
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Example
Calculate ∆Hrxn for the formation of CO,
given the reactions below:
– C(s) + 1/2 O2 (g) CO(g) ∆Hrxn = ? – C(s) + O2 (g) CO2(g) ∆Hrxn = -393.5 kJ
– CO(g) + 1/2 O2 (g) CO2(g) ∆Hrxn = -283.0 kJ
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Enthalpies of Formation
∆Hf is enthalpy change for formation of
compound from constituent elements
– Depends upon conditions of reactants andproducts
• Temp, pressure, state
– For simplicity, “standard state” is defined
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Standard State
Set of conditions (the ones found in
most laboratories) for which enthalpies
are recorded – Pure form of a substance
– Pressure = 1 atm
– Temperature = 298K (25ºC)
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∆Hº = Enthalpy at Standard
StateStandard enthalpy change
– Indicates that all reactants and products
are in standard state conditions – Can be for formation or reaction
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∆Hºf
Standard enthalpy of formation
–Change in enthalpy for reaction that forms
one mole of compound from elements – All substances in standard state
2 C(graphite) + 3 H2 (g) + 1/2 O2 (g) C2H5OH(l) ∆Hºf = -277.7 kJ
Note that all substances are in their most stable form
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Table of ∆Hºf Values
QuickTime™ and a
TIFF (Uncompressed) decompressor are needed t o see this picture.
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Question
True or False: to use Hess’s Law to
find the enthalpy of reaction, the
reaction MUST have taken place via aseries of steps, for which the
enthalpies of reaction are known.
a. TRUEb. FALSE
Careful! Tricky question!
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Question
For which of the following is the
enthalpy of formation ZERO?
a. O2(g)
b. O2(l)
c. O(g)
d. All of these
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Question
True or False: Because ethane (C2H6)
is most stable as a gas, the enthalpy
of formation of gaseous ethane isZERO
a. TRUE
b. FALSE
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Example
What is the equation for formation of
glucose in terms of standard enthalpy of
formation?
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Enthalpy of Reaction
Enthalpies of formation (∆Hºf ) can be
used to calculate standard enthalpy of
reaction (∆Hºrxn) – ∆Hºrxn = (sum of ∆Hºf of products) - (sum of
∆Hºf of reactants)
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Finding ∆Hºf
Find ∆Hºf for C3H6O(l), given the
equation:
C3H6O(l) + 4 O2(g) → 3 CO2(g) + 3 H2O(l), ∆H°
rxn = -1790 kJ
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Food and Fuel
Combustion is the burning of a chemicalto release heat and produce simpler
substancesFoods are burned in the presence ofoxygen to produce various products
– CO2
– H2O
– Chemical energy or heat
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Food Calories
A food Calorie is actually a kcal (1000
cal)
Three types of nutrients – Carbohydrates, 4 Cal/g
– Proteins, 4 Cal/g
– Fats, 9 Cal/g
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Heat IS Energy
The Calories in a meal can be used to
fuel biological processes
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Example
A meal consisting of a Double Quarter-
Pounder with Cheese, large fries, and large
Triple-Thick Chocolate Shake contains 2470
food calories. If the energy from this meal
were used to heat 10 gallons (approximately
38L or 38kg) of water initially at room
temperature (25°
C), how hot would thewater get?