Transcript 2 Thermo NOTES cp2 u7 1112

Question of the Day
1. A 1.0 mole sample of ethanol
(C2H5OH) has a heat capacity of
110.4 J/°C. Calculate the specific
heat of ethanol (hint: What is the
relationship between heat capacity
and specific heat?)
Day 6 2-22
= 2.4 J/g°C
Thermodynamics
Energy
Enthalpy
Entropy
Spontaneity
2
State Functions
• Property determined under
specified conditions of T, P,
location, etc
• Independent of HOW conditions are
reached
3
State Functions
Examples
Energy
d
B
B
_____________________________________
DE = difference
between A & B
_____________________________________
A
A
4
State Functions:
Hess’s Law
• For a reaction that occurs in
several steps, the sum of the
D’s for each step is the D for the
net reaction.
• Thus, it is possible to calculate
D for a complex reaction by
adding the D’s of each simple
reaction that leads to the
complete reaction.
5
State Functions:
Hess’s Law
• The Equation
DTOTAL = Dstep 1 + Dstep 2 + Dstep 3 + … + Dstep n
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Energy Transfers:
• Possibilities:
• From system to surroundings
• From surroundings to system
• How heat transfer is made:
• Heat = q
• Positive change = heat added to system
• Negative change = heat released from system
• Work = w
• Positive change = work done ON system
• Negative change = work done BY system
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Energy Transfers
• First Law of Thermodynamics
• Energy lost by the system
EQUALS energy gained by the
surroundings
• Energy gained by the system EQUALS
energy lost by the surroundings
8
Energy Transfers
• Exothermic Change
system
• Energy moves from ______________
to surroundings
________________
Loss of q
• _________
• Temperature of surroundings
increases
________________
9
Energy Transfers
• Endothermic Change
surroundings
• Energy moves from ________________
system
to _____________
Gain of q
• _________
• Temperature of surroundings
decreases
________________
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Energy Transfer
Equations
• Internal Energy: DE = Ef – Ei
• Ef = final energy
• Ei = initial/starting energy
• Energy Change due to Heat &
Work: DE = q + w
•
•
•
•
Use
Use
Use
Use
+ q if endo
– q if exo
+ w if work is done ON system
– w if work is done BY system
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Assignment
Review section 17.1 and
complete #s 5-11 on page 561.
- Due Monday 2-22
Assignment
Review section 17.1 and
complete #s 42, 43, 45, 47,
48, and 50-52 on page 586.
- Due Thursday 2-25
Question of the Day
1. The specific heat of ice is 2.1
J/g°C. What is the heat capacity
for a 2.0 mole sample?
Day 1 2-23
= 75.6 J/°C
Enthalpy
• Thermodynamic property that
measures/describes the changes in
heat content of a system under
constant pressure.
• Symbol: H, DH
• Units of measure: cal, kcal, J, kJ
• State function
• Equations and usage DH = Hf - Hi
+DH
Hf > Hi
Heat gained & endo
-DH
Hf < Hi
Heat lost & exo15
BOOK - Calorimetry
•Constant-Pressure Calorimeters
The enthalpy (H) of a system accounts
for the heat flow of the system at
constant pressure.
• The heat absorbed or released by a reaction
at constant pressure is the same as the
change in enthalpy, symbolized as ΔH.
BOOK - Calorimetry
•Constant-Pressure Calorimeters
The value of ΔH of a reaction can be
determined by measuring the heat flow
of the reaction at constant pressure.
• In this textbook, the terms heat and
enthalpy change are used
interchangeably.
• In other words, q = ΔH.
Calculating Enthalpy
Change in Reactions
• Formula:
DHRXN = DHPRODUCTS - DHREACTANTS
2 H2(g) + O2(g)
Higher E means
__________________________
Less stability
+
483.6
–
Lower E means
__________________________
more stability
2 H2O(g)
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Calculating Enthalpy
Change in Reactions
CH4(g) + 2O2(g)
+
H
CH4(g) + 2O2(g)
802 –
CO2(g) + 2H2O(g)
+
890
–
CO2(g) + 2H2O(l)
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Calculating Enthalpy
Change in Reactions
•Phase matters!!!!!!!
•Heat In = Heat Out
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Thermochemical
Equations
•Thermochemical Equations
How can you express the enthalpy
change for a reaction in a chemical
equation?
In a chemical equation, the enthalpy
change for the reaction can be written
as either a reactant or a product.
Thermochemical Equations
In the equation describing the exo reaction
of calcium oxide and water, the enthalpy
change can be considered a product.
CaO(s) + H2O(l) → Ca(OH)2(s) + 65.2 kJ
Calcium oxide is
one of the
components of
cement.
Thermochemical Equations
A chemical equation that includes the enthalpy
change is called a thermochemical equation.
CaO(s) + H2O(l) → Ca(OH)2(s) + 65.2 kJ
Thermochemical Equations
•Heats of Reaction
The heat of rxn is the ΔH for the chemical
equation exactly as it is written.
• Heats of reaction are reported as ΔH.
• The physical state of the reactants and
products must also be given.
• The standard conditions are that the
reaction is carried out at 101.3 kPa (1 atm)
and 25°C.
Thermochemical Equations
•Heats of Reaction
Each mole of calcium oxide and water that
reacts to form calcium hydroxide
produces 65.2 kJ of heat.
CaO(s) + H2O(l) → Ca(OH)2(s)
ΔH = –65.2 kJ
• In exothermic processes, the chemical
potential energy of the reactants is higher
than the chemical potential energy of the
products.
Thermochemical Equations
•Heats of Reaction
Baking soda decomposes when it is heated.
This process is endothermic.
2NaHCO3(s) + 85 kJ → Na2CO3(s) + H2O(l) + CO2(g)
The carbon
dioxide
released in the
reaction
causes muffins
to rise while
baking.
Thermochemical Equations
•Heats of Reaction
2NaHCO3(s) + 85 kJ → Na2CO3(s) + H2O(l) + CO2(g)
Remember that ΔH is positive for
endothermic reactions. Therefore, you can
write the reaction as follows:
2NaHCO3(s) → Na2CO3(s) + H2O(l) + CO2(g)
ΔH = 85 kJ
Thermochemical Equations
•Heats of Reaction
The amount of heat released or absorbed
during a rxn depends on the number of
moles of the reactant involved.
• The decomposition of 2 mol of sodium
bicarbonate requires 85 kJ of heat.
• Therefore, the decomposition of 4 mol of the
same substance would require twice as
much heat, or 170 kJ.
Thermochemical Equations
•Heats of Reaction
To see why the physical state of the reactants
and products must be stated, compare the
following two equations.
H2O(l) → H2(g) +
H2O(g) → H2(g) +
O2(g)
ΔH = 285.8 kJ
1
2
ΔH = 241.8 kJ
1
2
O2(g)
difference = 44.0 kJ
Sample Problem 17.4
Using the Heat of Reaction to Calculate
Enthalpy Change
2NaHCO3(s) + 85 kJ → Na2CO3(s) + H2O(l) + CO2(g)
•Calculate the amount of heat (in
kJ) required to decompose 2.24
mol NaHCO3(s).
Assignment
Review section 17.1 and
complete #s 42, 43, 45, 47,
48, and 50-52 on page 586.
- Due Thursday 2-25
Assignment:
#s 14 and 15 on page 567 (show
work), #s 16, 17, 20, and 21 on
page 568 (show work)
- Due Wednesday 2-17
Question of the Day
1. The appropriate unit for
specific heat =
2. Ice subliming (solid  gas) is
exo or endo?
Day 5 2-13
PRACTICE:
2NaHCO3(s) + 85 kJ → Na2CO3(s) + H2O(l) + CO2(g)
1. Calculate the amount of heat (in
kJ) required to produce 3 mols of
Na2CO3(s).
= 255 KJ
1. If a sample of NaHCO3(s) requires
127.5 kJ to completely decompose,
what was the mass of the original
sample?
= 252 g
Question of the Day:
3C2H2(g)  C6H6(l) + 630 kJ
Is this reaction endo OR exo?
How much heat is involved if you
start with 2 grams of acetylene
(acetylene = C2H2)?
Day 6 217
EXO 16.2 kJ
Thermochemical
Equations
•Heats of Combustion
The heat of combustion is the heat of
reaction for the complete burning of one mole
of a substance.
Thermochemical Equations
•Heats of Combustion
Small amounts of natural gas within crude
oil are burned off at oil refineries.
CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) + 890 kJ
• This is an exothermic reaction.
• Burning 1 mol of methane
releases 890 kJ of heat.
• The heat of combustion (ΔH)
for this reaction is –890 kJ per
mole of methane burned.
Heats of Combustion at 25°C
Interpret Data
ΔH (kJ/mol)
Substance
Formula
Hydrogen
H2(g)
–286
Carbon
C(s,
graphite)
–394
Methane
CH4(g)
–890
Acetylene
C2H2(g)
–1300
Ethanol
C2H6O(l)
–1368
Propane
C3H8(g)
–2220
Glucose
C6H12O6(s)
–2808
Octane
C8H18(l)
–5471
Sucrose
C12H22O11(s)
–5645
Like other heats of
reaction, heats of
combustion are
reported as the
enthalpy changes
when the reactions
are carried out at
101.3 kPa and
25°C.
•Which of the following
thermochemical equations represents
an endothermic reaction?
A. Cgraphite(s) + 2 kJ
B. 2H2(g) + O2(g)
Cdiamond(s)
2H2O + 483.6 kJ
Vocab. from the book:
Heats of reaction
Heats of combustion
Molar heat of fusion (melting)
Molar heat of solidification
(freezing)
ΔHfus = -ΔHsolid
Vocab. from the book:
Molar heat of vaporization
Molar heat of condensation
ΔHvap = -ΔHcond
Molar heat of solution (dissolving)
Spontaneity
• Main question of chemistry:
Will a reaction GO! when reactants are put
together?
• Spontaneous reactions
• Reaction occurs by itself, activation E may be needed
• May have fast, moderate, or slow rate
• Non-spontaneous reactions
• Require constant E supply to occur
• Reaction stops when E supply is removed
• Reversing reactions
• If a reaction is spontaneous under specified
conditions, the reverse is non-spontaneous at those
same conditions
42
Predicting Sponaneity
• Why is DH a pretty good
predictor of spontaneity?
• -DH : heat out
• +DH: heat supply needed
• When might DH not indicate
spontaneity well?
43
Predicting Spontaneity
• H2O(s) → H2O(l)
∆H = + 6.0 kJ;
1 atm, 0°C
44
Predicting Spontaneity
• H2O(l) → H2O(g)
∆H = + 40.7 kJ;
1 atm, 100°C
45
Predicting Spontaneity
• CaCO3(s) → CaO(s) + CO2(g)
∆H = +178.0 kJ;
298 K vs. 1100 K
46
Predicting Spontaneity
47
Question of the Day:
Day 1
2-18
NaCl(s) + 3.88kJ  NaCl(aq)
Is this reaction endo OR exo?
Draw an enthalpy diagram
representing the change above:
NaCl(aq)
+
3.88 kJ
NaCl(s)
–
Entropy
• Entropy is the measure of the
or randomness of a system.
disorder
• Entropy is a state function.
• Entropy changes follow Hess’ Law.
• A system has high entropy if it
• _____________________________________________
(Very disorganized)
• _______________________________________
Is Very dsorigniaezd
Has freedom of motion
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Entropy & Spontaneity
• Reactions will be spontaneous
etnopry
if
________ is increased.
• Examples
• ice cubes melting vs. putting
water in freezer to make ice cubes
• room gets messy vs. doing work to
put things in order
50
Standards & Constants
• Conditions & Units
• T = 25oC = 298K
• P = 1 atm
• Units: J/K = Joules/Kelvin
• What is the relationship between
entropy and temperature?
• As T 0 K, S  0
• T↓, S↓ and T↑, S↑
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Standards & Constants
• Why are the entropy values for
elements and compounds always
positive?
• T always > 0K
• Therefore everything is always moving
• How do entropies compare among
the phases of matter?
• So
solid
< So
liquid
< So
gas
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Standard Entropy
Changes
• The equation
∆So =
SSoproducts – SSoreactants
• Example:
2HCI(aq) + 2Ag(s) → 2AgCl(s) + H2(g)1 atm, 25oC
• Predict if this reaction will have + ∆So or – ∆So.
• Calculate ∆So (NEED STANDARD VALUE
HANDOUTS).
53
Standard Entropy
Changes
• The equation
DSo =
SnSoproducts – SmSoreactants
• Example:
2 H2(g) + O2(g) → 2 H2O(g)
1 atm, 25oC
• Predict if this reaction will have + ∆So or – ∆So.
• Calculate ∆So.
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Quick Talk:
Enthalpy
ΔS
q
State function
Entropy
ΔH
Joules
4.184 J
Potential Chemical
Specific
Energy
heat
Assignment:
#s 14 and 15 on page 567 (show
work), #s 16, 17, 20, and 21 on
page 568 (show work)
- Due Thursday 2-19 (pd 1)
Wednesday 2-18 (pd 3)