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Unit 7:
(Chapter 6)
Thermochemistry
I. Energy
A. Energy is the capacity to do ____________. It’s two major forms:
1. Potential energy: Due to its _______________ or due to forces of attraction or repulsion.
a. The position of an object in a gravitational field determines its PE.
b. The position of a charged object in an electric field determines its PE.
c. Chemical energy (due to the _______________ of chemical bonds) can be considered PE.
The relative positions of the unbonded atoms versus the bonded atoms results in an
energy change. Example: H + H  H2
2. Kinetic energy: Energy of ____________. The motion of the molecules can be translational, but also
from vibrations and rotations.
B. Units of Energy
1. The SI unit of energy is the joule (J):
J
kg m 2
or 1 kJ = 1000 J.
s2
2. The calorie is defined as the heat needed to raise the temperature of ______ of water by _____.
a. 1 cal = ___________
 J
b. 1000 cal = 1 kcal = 1 Calorie (food calorie)
II. Thermochemistry: Basic Terms
A. Thermochemistry considers the energy changes that occur during physical and ______________ changes.
B. System: Portion of the universe we are measuring.
1. Surroundings: The rest of the universe (generally only consider part that is nearby, that which will
have an effect on the system).
2. Open system: Matter and energy can exchange freely with the surroundings.
3. Closed system: Only energy can exchange freely with the surroundings (not matter).
4. Isolated system: Does not exchange energy or matter with the surroundings.
C. Heat (q) is a transfer of energy caused by a ___________________ difference. When no heat flows between
the system and the surroundings, the two are in thermal __________________ or at the same temperature.
1. Exothermic Process (-): System _________________ heat to the surroundings (heat exits).
Example: _________________ reactions
2. Endothermic Process (+): System __________________ heat from the surroundings (heat enters).
Example: Ice melting
D. Work (w) is a transfer of energy caused by a force applied through a distance.
III. The First Law of Thermodynamics
A. Energy cannot be created nor destroyed in a chemical or physical process.
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B. Internal energy (U) of a system is the total energy within the system (PE and KE).
C. The change in internal energy can be calculated by the following equation:
U  q  w
D. Internal energy is a state function: it does not depend on the path.
See diagram to the right.
E. Heat (q) and work (w) are path functions: they do depend on the path.
F. In an isolated system, the change in internal energy is __________.

G. In a closed system, if the system gains energy as heat, but loses an
equal quantity of energy as work, the ∆U = ____________.
H. Heat (q) is the easiest of these quantities to measure experimentally.
IV. Enthalpy Change (∆H) and Chemical Reactions
A. ALL chemical reactions have heat changes: _________________ or ________________.
1. Exothermic process: Heat is _________________ . Chemical energy is converted to thermal energy
and is transferred to the __________________. (__________________)
Reactants → Products + ________________
(a) The heat was lost from the system to the surroundings instantaneously.
(b) The heat is lost to the surroundings slowly, so the system temperature increases.
(c) In an isolated system, essentially no heat is lost to the surroundings, so the temperature
rises in the system to the greatest extent possible.
2. Endothermic process: Heat is ___________________ by the reaction. Thermal
energy is transferred from the surroundings to the _____________. (______________)
Reactants + ________________→ Products
B. Heat of reaction (q) can be determined at constant volume (qv) or constant pressure (qp). Most chemical
reactions, though, are carried out at constant pressure. Because of this the term enthalpy (H) was
defined as qp.
C. Enthalpy
1. Extensive property: depends on the quantities of substances present. If 2 mol of methane is burned
it will give off twice as much heat as compared to if 1 mol is burned.
2. State function: does not depend on the path.
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EXOTHERMIC RXN → ΔH = _______________
ENDOTHERMIC RXN → ΔH = _______________
3. Forms of ΔH:
ΔHrxn
ΔHc
ΔHf
ΔHsoln
enthalpy of reaction
enthalpy of combustion
enthalpy of formation
enthalpy of solution
(kJ)
(kJ / mol of combustible substance)
(kJ / mol of substance formed)
(kJ / mol of solute)
4. The combustion of methane: ΔHc = - 890.3 kJ/mol
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(l) + 890.3 kJ
5. The decomposition of mercuric oxide:
2 HgO(s)  2 Hg(l) + O2(g)
HgO(s)  Hg(l) + ½ O2(g)
ΔHrxn = +181.66 kJ
ΔHrxn = _________
6. The synthesis of hydrogen iodide:
H2(g) + I2(s)  2 HI(g)
ΔHrxn = +52.96 kJ
HI(g)  ½ H2(g) + ½ I2(s)
ΔHrxn = _________
7. What is the enthalpy change if 125 g hydrogen iodide is decomposed?
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D. Techniques to determine the ΔH of a reaction:
1. Calculate it using average bond energies.
2. Measure if with a calorimeter.
3. Calculate it using Heats of Formation data. (Appendix C)
4. Calculate it using Hess’ Law.
V. Estimation of ΔH Using Average Bond Energies (Section 9.10, p. 366 - 371)
A. Bond energy: The energy required to ___________ a chemical bond. Energy is ALWAYS
____________, therefore the process is ______________ and the value is _____________.
Bond-dissociation energy (D) is the quantity of energy required to break 1 mol of covalent bonds between
two atoms in a molecule in the gas phase.
1. Measurement of bond energy for diatomic molecules is relatively easy:
Bond Breaking: H2(g) → H(g) + H(g)
Bond Forming:
H(g) + H(g) → H2(g)
ΔH = +436 kJ/mol (ENDO)
ΔH = - 436 kJ/mol (EXO)
2. Determining the bond energy for a polyatomic molecule is more complicated. Once one bond is
______________, the chemical environment changes, so it changes the amount of energy needed to
break a bond.
Example:
(a) H2O(g) → H(g) + OH(g)
(b) OH(g) → H(g) + O(g)
ΔH = +499 kJ/mol
ΔH = +428 kJ/mol
3. The bond energy of an O-H bond changes in different molecules, therefore the_____________ bond
energy is determined and used in calculations (____________kJ/mol).
4. Table of Average Bond Energies (Table 9.1, p. 367)
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5. Bond formation is always an _________________ process, therefore a _____ sign must be added to
the data when referring to bond formation.
Example:
F(g) + F(g) → F2(g)
ΔH = __________
6. Bond energy increases as the bond length _____________. Therefore smaller atoms will have a
____________ bond than larger atoms.
Example:
H–H
________
Cl – Cl
________
7. Double and triple bonds are ______________ than single bonds.
Example:
N≡N
________
N–N
________
B. Estimating ΔH from average bond energies
1. Write the balanced chemical equation.
2. Draw Lewis structures for each substance.
3. Tabulate the bonds broken (reactants) and bonds formed (products) and find the E of each bond.
4. Determine ΔH using the following equation:
ΔH = Σ E reactants + Σ E products
(+)
(-)
Negative sign MUST be added for the products
OR
ΔHrxn = Σ E reactants – Σ E products
C. Examples:
1. Estimate the ΔHc of methane using average bond energies.
2. Estimate the ΔHc of hydrogen gas.
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3. Estimate the enthalpy of formation of gaseous hydrazine (N2H4).
4. Estimate the ΔHrxn of C2H6(g) + Cl2(g)  C2H5Cl(g) + HCl(g).
VI. Calorimetry: Measuring Quantities of Heat
A. Calorimeter: An instrument that measures the heat content.
1. Simple calorimeter:
2. Bomb calorimeter (ΔHc)
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B. The relationship between heat (q) and temperature:
1. Heat capacity (C): The quantity of heat needed to raise the temperature of an object by 1°C.
(Generally used for objects like bomb calorimeters)
a. Calculations with heat capacity:
q = CΔT
2. Specific heat capacity (c): The quantity of heat needed to raise the temperature of 1 g of a
substance by 1°C.
a. Every substance has a unique specific heat. ______________
has an extremely high specific heat, meaning that it takes a lot of
heat to raise the temperature of 1 g of water by 1°C.
b. Specific heat of water in calories:______________
c. Which has a higher specific heat: sand or water?
d. Calculations with specific heat:
q = mcΔT
3. Molar heat capacity: The quantity of heat needed to raise the temperature of 1 mol of a substance
by 1°C.
C. Calculations with specific heat and heat capacity:
1. Calculate the heat capacity of an aluminum block that must absorb 629 J of heat from its
surroundings to raise its temperature from 22°C to 145°C? Given aluminum’s specific heat, what is the
mass of the Al block?
2. How much heat (in kJ) must be added to change the temperature of 1.00 gallon of water from 20.0°C
(room temperature) to 100.0°C? How many Calories would this be?
3. What is the final temperature of 100.0 mL of water at 80.0°C if it releases 55.0 kcal of heat to the
surroundings?
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D. Measuring Specific Heats: The specific heat of an object can be found by heating an
object and measuring how much heat is lost when it is placed in water of known temperature.
1. Basic Law of Heat Exchange:
Heat lost by hot substance = Heat gained by water
2.
qH2O = -qobject
 mcΔT = -(mcΔT)  mc(Tf-Ti) = -mc(Tf-Ti)
OR
mHcH(TH – TF) = mCcC(TF – TC)
3. Example Calculations:
a. Place a 150.0g block of aluminum into a pan of boiling water (T H = 100.0°C). Once the block has equilibrated with the
water, place it into 200.0 mL of 10.0°C water. As the block cools, the water warms up. The final temperature of the
mixture is 21.7°C. Determine the specific heat of Al and the % error for the experimental value. (c Al accepted = 0.900
J/g°C).
b. A 15.5 g sample of a metal alloy is heated to 98.9°C and then dropped into 25.0 g of water in a calorimeter. The
temperature of the water rises from 22.5 to 25.7°C. Calculate the specific heat of the alloy.
c. If a student mixes 475 mL of 95.0°C water and 225mL of 42.0°C water, what will the final temperature be just after
mixing?
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d. Many people believe that placing a cold spoon in a cup of hot coffee will cool it enough to drink comfortably. A 100.0 g
silver spoon at 15.0°C is placed in a 200.0 mL cup of coffee (4.10 J/g°C) at 97.0°C. What is the final temperature of the
mixture?
E. Heats of Reaction and Calorimetry (NOTE: For all calorimetry problems, solve for q using a positive
ΔT, then add a negative if the reaction / process is exothermic!)
1. Examples with simple calorimeters (Styrofoam cup):
a. Two solutions are mixed in the calorimeter: 40.0 mL of 1.00 M KOH and 40.0 mL of 0.500 M H 2SO4, both at 21.00°C.
The final temperature has a density of 1.02 g/mL, a volume of 80.0 mL, a specific heat of 4.00 J/g°C, and a temperature of
27.85°C. Calculate the ΔH per mol of water formed.
b. Calculate the ΔHsoln when 5.00 g of a solute (molar mass = 120 g/mol) is dissolved in 150.0 mL of water if the
temperature is lowered by 4.80°C. Determine the value in kJ/mol and kcal/g.
2. Examples with bomb calorimeters (reactions that involve gases, which must be done in a
closed container; most often used for combustion reactions).

To use a bomb calorimeter, we must know the HEAT CAPACITY (C) by measuring the
temperature change when burning a substance of known ΔHc.

Benzoic acid is often used to calibrate the bomb calorimeter:
C6H6COOH(s) + 15/2 O2(g)  7 CO2(g) + 3 H2O(l) ΔHc = -26.42 kJ/g
a. The combustion of 1.045 g of benzoic acid caused the temperature of a bomb calorimeter to rise from 21.02°C to
25.23°C. What is the heat capacity (C) of the calorimeter? (6.56 kJ/°C)
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b. When the SAME CALORIMETER as in part (a) burns 0.946 g ethanol, the temperature changes from 20.65°C to
24.92°C. Determine the heat of combustion of ethanol. (-1.36 x 103 kJ/mol)
C2H5OH(l) + 3 O2(g)  2 CO2(g) + 3 H2O(l)
*** NOTE: The above calculation is not the enthalpy change for the reaction since the bomb calorimeter is at
constant VOLUME and not constant pressure. It is really a measure of internal energy, but the difference
between the two is so small we can assume they are equal.
ΔH = ΔU + PΔV = ΔU - ΔngasRT
For the ethanol reaction: (-1 mol)(8.314 J/mol K)(297.92 K) = -2.4 kJ (negligible)
c. The heat capacity of the bomb calorimeter is found to be 5.15 J//°C. If that calorimeter is used to combust 0.480 g of
graphite with an excess of oxygen and the temperature rises from 25.00°C to 28.05°C, calculate the heat of combustion of
graphite. (-393 kJ/mol)
VII. Hess’ Law of Heat Summation
A. Hess’ Law: The enthalpy change of a reaction is the same whether it is run in one step or several.
What is the ΔH(a)?
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B. Sometimes a compound is difficult to synthesize from its elements, so the direct measurement of ΔH f cannot
be done. In these cases, we can measure ΔHf using Hess’ Law.
C. When reactants are converted to products, the ΔH is the same whether the reaction takes place in ______
step or multiple _________. For example: if you go to from the first to the sixth floor by elevator or by stairs, your
net gain in potential energy is the same.
Equation:
ΔHrxn = ΔH1 + ΔH2 + ΔH3 + …
D. Steps for Calculation Using Hess’ Law:
1. Write the net reaction.
2. Write the given reactions (if not already done), including ΔHrxn (in kJ)
3. Rearrange the given reactions so that they sum to the net reaction.
4. If the direction of a reaction is reversed, the ΔH has the opposite sign.
5. If you multiply the coefficients of the reaction by a factor, the ΔH is multiplied by the same factor.
6. Sum the values of ΔH.
7. Calculate ΔHf or ΔHc if requested.
E. Examples:
1. Calculate the ΔHrxn of NO + O  NO2
Given: NO + O3  NO2 + O2
O3  3/2 O2
O2  2 O
(ANSWER = -306 kJ)
ΔH = -200 kJ
ΔH = -143 kJ
ΔH = 498 kJ
2. Calculate the ΔHf of carbon monoxide gas when it is formed from graphite and oxygen gas.
Given: ΔHf CO2(g) = –394 kJ/mol
ΔHc CO(g) = –283 kJ/mol (CO2 is formed)
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3. Calculate the ΔHf of acetylene (C2H2).
Given: ΔHf CO2(g) from graphite = -393.5 kJ/mol
ΔHf H2O(l) = -285.5 kJ/mol
ΔHc C2H2(g) = -1299.6 kJ/mol
VII. Standard Enthalpies of Formation and Reactions
A. Standard heat (enthalpy) of formation: The energy ____________ or ___________ when one mol of a
substance is formed from its _____________ at _______ and 1 atm.
B. Symbol:
ΔH°f
C. Data is found in Appendix C (and attached)
D. Elements have ΔH°f values of ______ by definition at standard state (25°C, 1 atm).
E. Examples: Values depend on the state of matter!!!
Br2(l)
Br2(g)
Br(g)
Br-(aq)
H2O(l)
CH4(g)
F. The calculation of ΔHrxn can be determine directly using the heats of formation data and
the following equation:
ΔHrxn = Σ ΔH°f products - Σ ΔH°f reactants
**NOTE: ΔHrxn has units of kJ **
G. Examples:
1. Calculate the ΔHrxn and ΔHc for the combustion of methane gas (CH4).
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2. Determine the enthalpy of combustion of ethanol.
VIII. Combustion and Respiration: Foods and Fuels
Fuel Value: Energy ______________ when a fuel is combusted (ΔH c). Because it is defined as the energy released,
the number is always ______________. For foods, it is often called ____________
value.
Nutritional Calorie = 1 Cal = 1 kcal = 1000 cal
A. Fuel: A substance that burns with the release of heat.
1. Some common fuels:
Fuel
wood (pine)
anthracite coal
crude oil
gasoline
natural gas
hydrogen
Fuel Value (kJ/g)
18
31
45
48
49
142
2. We are using fossil fuels 50,000 times faster than they were formed, therefore they WILL be depleted
unless we change to renewable sources (such as: ____________________________________).
3. Coal is a complex organic material with ______________ as the primary element. It also contains
H, O, N, and S. Complete combustion of carbon produces__________.
a. Burning coal also produces CO and soot (unburned carbon). Noncombustible matter is left
unburned as ashes (fly ash). Sulfur dioxide is produced (which oxidizes to ______________ in
the atmosphere). Nitrogen oxides are also produced (because at high T, nitrogen and oxygen
in air combine to form nitrogen monoxide).
4. Natural gas is mainly ____________, with some ____________, ____________, and ___________.
Natural gas gives almost twice the energy as coal. Combustion of natural gas does produce CO, soot,
nitrogen oxides, but no ash or sulfur dioxide.
5. Petroleum is a complex mixture of hydrocarbons. Petroleum products (like gasoline) are cleaner
burning than coal (but not as clean as natural gas). It is a mixture of hydrocarbons with 5 to 12 carbon
atoms. Octane (__________) is used as a representative molecule.
a. Difference between gasoline of different octane ratings?
Scale: 0 = pure heptane (burns explosively, leads to knocking sound)
100 = pure isooctane (smooth burning)
b. All octane ratings give the SAME AMOUNT OF ENERGY! So mpg does not improve by
using the higher octane gas. Most cars work best on 87-octane.
B. Foods: Fuel for the body.
1. Three principal classes of foods: carbohydrates, fats, and protein.
Average nutritional values:
Fat
Protein
Carbohydrate
9.00 kcal / g
4.00 kcal / g
4.00 kcal / g
38 kJ/g
17 kJ/g
17 kJ/g
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2. Bomb calorimetry produces the same caloric value as do metabolic studies for starches, sugars,
and fats (Hess’ Law).
3. Fats yield relatively higher energies on combustion than do carbohydrates. This means they are
efficient chemical store of energy for the body (more energy stored per mass). The body can store
up to about 500 g of carbohydrate (as glycogen), but fat storage is unlimited!!
Combustion of glucose: C6H12O6(s) + 6 O2(g) → 6 CO2(g) + 6 H2O(l)
ΔHc = -2801 kJ/mol = -3.715 kcal/g
Nutritional value: +3.715 Cal/g
Combustion of a fat: 2 C57H110O6(s) + 163 O2(g) → 114 CO2(g) + 110 H2O(l)
ΔHc = -37,760 kJ/mol = -10.121 kcal/g Nutritional value: +10.121 Cal/g
4. Fats belong to a family of organic compounds called esters.
5. A bomb calorimeter (or a simple calorimeter) can be used to find the caloric content (or fuel value) of a
food.
ΔHc= Heat gained by H2O
mass of food burned
ΔHc= ___mc ΔT__
mass of food
a. When 2.50 g of a food is burned under a simple calorimeter that contains 200.0 mL of
15.0°C water, the temperature of the water rises to 53.8°C. How many Cal/g does this food
contain?
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b. (a) The human body is about 67% water by mass. What mass of sucrose (C 12H22O11,
ΔH°f = -2225 kJ/mol) must be metabolized by a 55 kg person with hypothermia to raise the
temperature of the body water from 33.5°C to the normal 37.0°C. Assume the products of
metabolism are at 25°C. (b) What volume of air at 37°C having a partial pressure of O2 of
151 Torr is required for this metabolism? (33 g, 1.5 x 102L)
IX. Thermodynamics (Chapter 17): The study of the interconversions of ________ and other forms of energy.
A. Entropy: A measure of the ___________________ or disorder of a system.
1. In nature, everything proceeds towards ___________ entropy and ______________enthalpy.
2. The greater the disorder, the _____________ the entropy of a system.
3. A gas has ________________ entropy than a solid.
4. The symbol for entropy is a capital ____ and has units _________(therefore it is dependent upon
_______________).
5. The probability of that a given substance will exist in a specified state is called_______________
entropy.
a. The symbol S° indicates that it is at standard conditions (____°C, _____atm).
b. As S° increases, the probability of existence ______________.
6. Entropy changes in reactions: represented by ________.
B. Calculation of ΔS
Equation:
ΔSrxn = Σ ΔS° products - Σ ΔS° reactants
1. ALL substances have an absolute entropy, including _______________.
examples:
2. Units for ΔSrxn = _______
3. ΔS is favorable when entropy is ________________ (ΔS is _________________).
4. ΔS is unfavorable when entropy is ________________( ΔS is negative).
5.
FAVORABLE
ΔH
ΔS
______
______
UNFAVORABLE
______
+ = increasing
______
- = decreasing
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6. Ex. Calculate the ΔS°rxn for the reaction:
Cl2(g) + 2 HBr(g) → 2 HCl(g) + Br2(g)
C. Spontaneous Processes
1. Spontaneous: The process WILL _________ without outside intervention. It is a
reaction that released free ___________ (usually as heat) and moves to a __________,
more thermodynamically stable state.
2. Spontaneous reactions can occur extremely ______________. As long as they will
occur, they are considered spontaneous.
Example: Oxidation of a diamond
3. To determine if a reaction is spontaneous, both driving forces must be determined:
Reactions tend towards ______________ enthalpy and _____________ entropy.
D. Gibbs’ Free Energy (G)
1. It’s a measure of both the ________________ and ______________ changes of a system.
2. It is the “free” energy available to do _________.
3. It can be used to determine the spontaneity of a reaction.
If ΔG is positive  the reaction is NOT spontaneous
If ΔG is negative  the reaction IS spontaneous
4. Equation to determine ΔG
ΔG = ΔH – T ΔS
ΔG = Gibbs’ Free Energy (kJ)
ΔH = Enthalpy (kJ)
ΔS = Entropy (kJ/K)
T = Temperature (K)
***Convert ΔS units to kJ/K***
5. Reactions (or processes) will occur spontaneously if ΔG is _________________.
Example: Calculate the ΔH, ΔS, and ΔG of this reaction. State whether the reaction is exothermic or endothermic,
whether the entropy is favorable or unfavorable, and whether the reaction is spontaneous or non-spontaneous.
2 KClO3(s)  2 KCl(s) + 3 O2(g)
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E. Temperatures at which processes are likely to occur
ΔH
ΔS
Reaction occurs spontaneously at _______ temperatures. It occurs
when _________>_________; therefore T> ΔH/ ΔS.
ΔG is ALWAYS ___________________, therefore the reaction is
______________________. It is always spontaneous in the __________
direction.
ΔG is ALWAYS ___________________, therefore the reaction is
______________________.
Reaction occurs spontaneously at _______ temperatures. It occurs
when _________<_________; therefore T< ΔH/ ΔS.
Example: Given ΔH and ΔS, calculate ΔG for this reaction. At what temperatures is this reaction spontaneous? ΔH =
847.6 kJ; ΔS = 41.2 J/K
Al2O3(s) + 2 Fe(s)  2 Al(s) + Fe2O3(s)
(at standard conditions)
F. Entropy changes in solution processes:
1. As a solid dissolves in a liquid, the entropy ______________ because the motion
of the particles increases, and there is increased randomness. Even though
ΔH is positive (______________________) for the process, it can still
dissolve due to the increase in entropy, which is _________________. The
__________ the temperature, the more soluble a solid becomes due to the
increased entropy of the system. (ΔH = +, ΔS = +; soluble at_________ temps)
Ex. Solubility of salt increases as the temperature _____________.
2. As a gas dissolves in a liquid, the entropy ________________ because the gas
has more disorder (it can move freely in all directions) than the aqueous form.
To increase the solubility of a gas in a liquid, the temperature must be
________________, to reduce the unfavorable entropy change.
(ΔH = -, ΔS = -; soluble at ______ temps)
Ex. Is CO2 more soluble in warm or cold soda? In which one do you see more
bubbles of CO2?
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G. Gibbs’ Free Energy of Formation ΔG°f
1. ΔG can also be calculated using the following equation:
ΔGrxn = Σ ΔG°f products - Σ ΔG°f reactants
2. Calculate the ΔGrxn for the combustion of ethane (C2H6).
3. CaCO3(s)  CaO(s) + CO2(g) (at standard conditions)
(a) Find ΔH, ΔS, and ΔG of the reaction.
(b) Is the reaction spontaneous?
(c) At what temperatures is this reaction spontaneous?
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