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CHEMISTRY - UTEXAS 1E

CH.6 - THERMODYNAMICS

CONCEPT: ENERGY CHANGES AND ENERGY CONSERVATION

_______________________ is the branch of physical science concerned with heat and its transformations to and from

other forms of energy.

_______________________ is the branch of chemistry that deals with the heat involved in chemical and physical changes.

Energy Changes and Energy Conservation

• The ______________________ is the specific part of the universe that we are focused on.

• The _________________________ deals with everything outside of it.

When talking about the movement of energy or heat between the ____________________ & ____________________ we

use the terms: Endothermic & Exothermic.

2 H2 (g) + O2 (g) 2 H2O (l) +

HEAT

HEAT

+ 2 HgO (s) 2 Hg(l) + O2 (g)

EXAMPLE: Classify each of the following process as either exothermic or endothermic: a) Fusion of Ice.

b) Sublimation of CO2.

c) Vaporization of aqueous water.

d) Deposition of chlorine gas.

e) Condensation of water vapor.



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CONCEPT: ENERGY FLOW TO AND FROM A SYSTEM

The _______ Law of Thermodynamics states that energy cannot be created nor destroyed, but only converted from one

form to another.

In chemistry, we are normally concerned with the energy changes associated with the system, not with its surroundings.

∆E = q + w q = ∆H (enthalpy) w = - P∆V

∆E =

q = * For q: (+) when system __________, __________, __________, heat or energy,

(-) when system __________, __________, __________, __________ heat or energy.

w = * For w : (+) when work done _____ system _____ the surroundings. Key word: volume ______________

(-) when work done _____ system _____ the surroundings. Key word: volume _______________

EXAMPLE: Which of the following signs on q and w represent a system that is doing work on the surroundings, as well as

losing heat to the surroundings?

a) q = - , w = - b) q = +, w = + c) q = -, w = + d) q = +, w = -



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PRACTICE: ENERGY FLOW TO AND FROM A SYSTEM

EXAMPLE: An unknown gas expands in a container increasing the volume from 4.3 L to 8.2 L at a constant pressure of 931

mmHg.

a. Calculate the work done (in kJ) by the gas as it expands. (1 L · atm = 101.3 J)

b. Using part A, calculate the internal energy of the system if the system absorbs 2.3 kJ of energy.

c. Using part B, calculate the internal energy of the system if the system does work against a vacuum.

PRACTICE: The reaction of nitrogen with hydrogen to make ammonia has an enthalpy, ∆H = - 92.2 kJ:

N2 (g) + 3 H2 (g) 2 NH3 (g)

What is in the internal energy of the system if the reaction is done at a constant pressure of 20.0 atm and the volume compresses from 10 L to 5 L?



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CONCEPT: CONSTANT-VOLUME CALORIMETRY

Every object has its own _________________________ (C), the quantity of heat required to change its temperature by 1 K.

C = qΔT

[in units of JK ]

_________________________________ (c), the quantity of heat required to change 1 gram of a substance by 1 degree K.

c = [in units of J

g•K ]

If we know c of a substance, we can algebraically solve the amount of heat absorbed or released:

q =

EXAMPLE: At constant volume, the heat of combustion of a particular compound is – 4621.0 kJ/mol. When 2.319 grams of

this compound (molar mass = 192.75 g/mol) was burned in a bomb calorimeter, the temperature of the calorimeter

(including its contents) rose by 3.138oC. What is the heat capacity of the calorimeter in J/K?



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PRACTICE: CONSTANT-VOLUME CALORIMETRY

EXAMPLE: In an experiment a 9.87 carat (1 carat = 0.200g) diamond is heated to 72.25oC and immersed in 22.08 g of

water in a calorimeter. If the initial temperature of the water was 31.0oC what is the final temperature of the water? (cdiamond =

0.519J

g• oC) (cwater = 4.184

Jg• oC

).

PRACTICE 1: A sample of copper absorbs 35.3 kJ of heat, which increases the temperature by 25oC, determine the mass

(in kg) of the copper sample if the specific heat capacity of copper is 0.385 J

g• oC.

PRACTICE 2: 50.00 g of heated metal ore is placed into an insulated beaker containing 822.5 g of water. Once the metal

heats up the final temperature of the water is 32.08oC. If the metal gains 14.55 kJ of energy, what is the initial temperature

of the water?



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CONCEPT: CONSTANT-PRESSURE CALORIMETRY The ______________ of a reaction can be calculated through the use of

a coffee-cup calorimeter.

EXAMPLE: You place 50.0 mL of 0.100 M NaOH into a coffee-cup calorimeter at 50.00oC and carefully add 75.0 mL of 0.100 M HCl,

also at 50.00oC. After stirring, the final temperature is 76.12oC. (Heat capacity and density of water: 4.184 J

g• oC and 1.00

gmL ).

HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)

a) Calculate qsoln (in J)

b) Calculate the enthalpy, ∆Hrxn (in J/mol), for the formation of water.



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CONCEPT: HEAT SUMMATION

Many reactions are difficult, even impossible, to carry out in a single chemical step.

• They may often times require multiple steps to get to the final products.

__________ Law states that the enthalpy change (∆H) of an overall process is the sum of the enthalpy changes of its

individual steps.

EXAMPLE: For the following example calculate the unknown ∆H from the given ∆H values of the other equations.

Calculate the ∆Hrxn for

S(s) + 32

O2 (g) SO3 (g) ∆H = ?

Given the following set of reactions:

12

S (s) + 12

O2 (g) 12

SO2 (g) ∆H1 = – 296.8 kJ

2 SO3 (g) 2 SO2(g) + O2 (g) ∆H2 = 198.4 kJ



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PRACTICE: HEAT SUMMATION

A. Calculate the ∆Hrxn for

CO(g) + NO (g) CO2 (g) + 12

N2 (g) ∆H = ?

Given the following set of reactions:

CO2 (g) CO (g) + 12

O2 (g) ∆H1 = 283.0 kJ

N2 (g) + O2 (g) 2 NO (g) ∆H2 = 180.6 kJ

B. Calculate the ∆Hrxn for

ClF (g) + F2 (g) ClF3 (g) ∆Hrxn = ?

Given the following reactions:

Cl2O (g) + F2O 2 ClF (g) + O2 (g) ∆Hrxn = - 167.4 kJ

4 ClF3 (g) + 4 O2 (g) 2 Cl2O (g) + 6 F2O (g) ∆Hrxn = 682.8 kJ

2 F2 (g) + O2 (g) 2 F2O (g) ∆Hrxn = -181.7 kJ



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CONCEPT: STANDARD HEATS OF FORMATION (∆HRXN)

In a _______________ equation, 1 mole of a compound forms from its elements. The ______________________________

(∆Hof) is the enthalpy change for the chemical equation when all the substances are in their standard states.

C (graphite) + 2 H2 (g) CH4 (g) ΔHfo = −74.9kJ

When calculating ∆Hof remember:

1) An element in its standard state (elemental state) is given an ∆Hof of zero.

Ex: Na (s) P4 (s) Cl2 (g) S8 (g)

2) Most compounds have a negative ∆Hof. 3) To find the ∆Hrxn use the following formula:

ΔHrxno = ΔHf(products)

o −ΔHf(reac tan ts)o

EXAMPLE: The oxidation of ammonia is given by the following reaction:

4 NH3 (g) + 5 O2 (g) 4 NO (g) + 6 H2O (g)

Calculate the ∆Horxn if the ΔHfo

value for NH3 , NO and H2O are – 45.9 kJ/mol, 90.3 kJ/mol and – 241.8 kJ/mol

respectively.

PRACTICE: Ibuprofen is used as an anti-inflammatory agent used to deal with pain and bring down fevers. If it has a

molecular formula of C13H18O2 determine the balanced chemical equation that would give you directly the enthalpy of

formation for ibuprofen.



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PRACTICE: STANDARD HEATS OF FORMATION (∆HRXN)

EXAMPLE: Use the following bond strength values (kJ/mol):

C–H 412 C–O 360 C=O 743 C–C 348 H–H 436

C=C 611 C≡C 837 C≡O 1072 O–H 464 O=O 498

Calculate the enthalpy of the reaction shown in the formula below:

H–C≡C–H + H–H + C=O

C

C C

O

HH

H H



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CONCEPT: THERMOCHEMICAL PROCESSES

______________________________ is the branch of physical science concerned with heat and its transformations to and

from other forms of energy.

In terms of a chemical reaction, you will learn that depending on certain conditions they can occur or not:

• A reaction that requires no outside energy source is classified as a natural process and is

______________________ .

• A reaction that requires a continuous energy source to happen is classified as an unnatural process and is

______________________ .

EXAMPLE 1: Which of the following statements is not true?

a) The reverse of a spontaneous reaction is always non-spontaneous.

b) A spontaneous reaction always moves towards equilibrium.

c) A highly spontaneous reaction can occur at a fast or slow rate.

d) It is possible to create a non-spontaneous reaction.

PRACTICE: Which of the following statements is/are true?

a) The rusting of iron by oxygen is a non-spontaneous reaction.

b) The addition of a catalyst to a reaction increases spontaneity.

c) The movement of heat from a cold object to a hot object is a non-spontaneous reaction.

d) The diffusion of perfume molecules from one side of a room to the other is a non-spontaneous reaction.

e) None of the above.



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CONCEPT: FIRST LAW OF THERMODYNAMICS Recall that our system represents the chemical reaction, while the surroundings represent everything else. In terms of systems there are 3 major types.

A(n) ____________ system involves the transferring of both matter and energy between system and surroundings.

A(n) ____________ system involves the transferring of neither matter and energy between system and surroundings.

A(n) ____________ system involves the transferring of only energy between system and surroundings.

The First Law of Thermodynamics states that energy cannot be created nor destroyed, but only converted from one form to another.

System SurroundingsHeat q

∆U = ∆E = q + w q = ∆H (enthalpy) w = - P∆V

∆U or ∆E = q = * For q: (+) when system __________, __________, __________, heat or energy.

(-) when system __________, __________, __________, __________ heat or energy.

w = * For w : (+) when work done _____ system _____ the surroundings. Key word: volume _________ .

(-) when work done _____ system _____ the surroundings. Key word: volume _________ .



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CONCEPT: ENTROPY AND SPONTANEOUS REACTIONS

The _______ Law of Thermodynamics states that molecular systems tend to move spontaneously to a state of maximum

randomness or disorder.

This disorder is also called entropy and uses the variable __________.

In general, as we move from a solid liquid gas then entropy will ___________________ and its

sign will be ___________________.

Conversely, if we move from a gas liquid solid then entropy will ___________________ and its

sign will be ___________________.

EXAMPLE 1: Which should have the highest molar entropy at 25oC?

a) Ga (l)

b) Ga (s)

c) Ga (g)

d) All of them have the same molar entropy.

EXAMPLE 2: Which substance has greater molar entropy.

a) CH4 (g) or CCl4 (l)

b) Ne (g) or Xe (g)

c) CH3OH (l) or C6H5OH (l)



Page 14

PRACTICE: ENTROPY AND SPONTANEOUS REACTIONS (CALCULATIONS 1)

EXAMPLE 1: Arrange the following substances in the order of increasing entropy at 25oC.

XeF4 (s) HI (g) BaO (s) H2 (g) Hg (l) Br2 (l)

EXAMPLE 2: Containers A and B have two different gases that are allowed to enter Container C. Based on the image of

Container C what is the sign of entropy, ∆So.

B CA

PRACITCE: An ideal gas is allowed to expanded at constant temperature. What are the signs of ∆H, ∆S & ∆G.

BA



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EXAMPLE: Consider the spontaneous fusion of ice at room temperature. For this process what are the signs for ∆H, ∆S,

and ∆G?

∆H ∆S ∆G

a) + + +

b) - + 0

c) - + -

d) + + -

e) - - -

PRACTICE: Consider the freezing of liquid water at 30oC. For this process what are the signs for ∆H, ∆S, and ∆G?

∆H ∆S ∆G

a) + - +

b) - + 0

c) - + -

d) - - +

e) - - -



Page 16


PRACTICE 1: Predict the sign of ∆S in the system for each of the following processes:

a) Ag+ (aq) + Br – (aq) AgBr (s)

b) CI2 (g) 2 CI – (g)

c) CaCO3 (s) CaO (s) + CO2 (g)

d) Pb (s) at 50oC Pb (s) at 70oC

PRACTICE 2: For each of the following reactions state the signs of ∆H (enthalpy) and ∆S (entropy):

a) Fusion of Ice.

b) Sublimation of CO2.

c) Vaporization of aqueous water.

d) Deposition of chlorine gas.

e) Condensation of a water vapor.



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CONCEPT: CALCULATING ENTROPY OF A SYSTEM

The 2nd Law of Thermodynamics states in terms of a system the entropy of a system increases spontaneously.

• Besides the system we also have our ___________________ and together they form the total ________________.

Thus, to calculate the total entropy change, ∆STotal, we use the following equation:

ΔSTotal =

So if,

∆STotal > 0, the reaction is ______________________

∆STotal < 0, the reaction is ______________________

∆STotal = 0, the reaction is ______________________

EXAMPLE 1: Calculate the standard entropy (in kJ) of reaction at 25oC for the following reaction:

N2 (g) + 3 H2(g) 2 NH3 (g)

The standard molar entropies of N2, H2 and NH3 are 191.5 , 130.6 and 192.3 respectively.



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PRACTICE: CALCULATING ENTROPY OF A SYSTEM (CALCULATIONS 1)

EXAMPLE: The oxidation of iron metal is given by the following reaction:

4 Fe (s) + 3 O2 (g) 2 Fe2O3 (s)

a) Calculate the ∆Ssystem if the standard molar entropies of Fe, O2 and Fe2O3 are 27.3 JK ⋅mol

, 205.0 JK ⋅mol

and 87.4

JK ⋅mol

respectively.

b) Calculate the ∆Ssurroundings if the reaction is spontaneous at 25oC. The standard molar enthalpy of Fe2O3 is – 824.2 kJmol

.

c) Calculate the ∆STotal and determine if the reaction is spontaneous or non-spontaneous under standard-state conditions?

PRACTICE: Diethyl ether (C4H10O2, MW = 90.1 g/mol) has a boiling point of 35.6oC and heat of vaporization of 26.7 kJ/mol. What is the change in entropy (in kJ/K) when 3.2 g of diethyl ether at 35.6oC vaporizes at its boiling point?



Page 19

PRACTICE: CALCULATING ENTROPY OF A SYSTEM (CALCULATIONS 2)

EXAMPLE 1: The normal boiling point of liquid propane is 231 K. What is the enthalpy of vaporization of liquid propane?



Page 20

CONCEPT: GIBBS FREE ENERGY

Chemists are generally interested in the system (the reaction mixture) rather than the surroundings. In order to define the free energy of a chemical system they use the following equations:

If ∆G < 0, the reaction is _________________________

If ∆G > 0, the reaction is _________________________

If ∆G = 0, the reaction is _________________________

EXAMPLE 1: Which of the following statements is true for the following reaction?

N2O4 (g) 2 NO2 (g) ∆Hº = - 57.1 kJ ∆Sº = 175.8 kJ

a) The reaction is spontaneous at all temperatures.

b) The reaction is spontaneous at low temperatures.

c) The reaction is spontaneous at high temperatures.

d) The reaction is non-spontaneous at all temperatures.

ΔGo = ΔHo −TΔSo ΔG = ΔGo +RTlnQ



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PRACTICE: GIBBS FREE ENERGY (CALCULATIONS 1)

EXAMPLE: The reduction of iron(III) oxide with hydrogen produces iron metal and can be written as follows:

Fe2O3 (s) + 3 H2 (g) 2 Fe (s) + 3 H2O (g) ∆Hº = 98.8 kJ ∆Sº = 141.5 Jk

a) Is this reaction spontaneous under standard-state conditions at 25ºC? If not, at what temperature will it become spontaneous?

PRACTICE 1: If ∆G is small and positive which of the following statements is true?

a) The forward reaction is spontaneous and the system is far from equilibrium.

b) The forward reaction is spontaneous and the system is near equilibrium.

c) The reverse reaction is spontaneous and the system is far from equilibrium.

d) The reverse reaction is spontaneous and the system is near equilibrium.

PRACTICE 2: Nitrogen gas combines with fluorine gas to form nitrogen trifluoride according to the reaction below at 25ºC:

N2 (g) + 3 F2 (g) 2 NF3 (g) ∆Hº = -249.0 kJ ∆Sº = -278 J/K

Calculate ∆Gº and state if the reaction favors reactants or products at standard conditions.

a) ∆Gº = - 332 kJ; the reaction favors the formation of reactants.

b) ∆Gº = - 166 kJ; the reaction favors the formation of products.

c) ∆Gº = - 166 kJ; the reaction favors the formation of reactants.

d) ∆Gº = - 332 kJ; the reaction favors the formation of products.



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EXAMPLE: For mercury, ∆Hvap = 58.5 kJmol

and ∆Svap = 92.9 Jk ⋅mol

at 25ºC. Does mercury boil at 350ºC and 1 atm

pressure?

EXAMPLE: The chemical reaction, 2 NO2Br (g) 2 NO2 (g) + Br2 (g), has a ΔSo = 135 Jmol ⋅k

and ΔHo=

926 kJmol

. Calculate the temperature when Keq = 4.50 x 105.



Page 23


EXAMPLE 1: Calculate ∆Grxn at 25ºC under the conditions shown below for the following reaction.

3 Cl2 (g) 2 Cl3 (g) ∆Gº = + 31.6 kJ

The partial pressures of Cl2 and Cl3 are 0.83 atm and 4.9 atm respectively.

EXAMPLE 2: For the reaction: N2 (g) + 2 O2 (g) 2 NO2 (g), ΔGo = 75, 550 Jmol

at 175 K and ΔGo= 41,875

Jmol

at 225 K.

Calculate ΔSo and ΔHofor the reaction.



Page 24

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CHEMISTRY - UTEXAS 1E CH.6 - THERMODYNAMICSlightcat-files.s3.amazonaws.com/packets/admin_chemistry...CHEMISTRY - UTEXAS 1E CH.6 - THERMODYNAMICS Page 2 CONCEPT: ENERGY FLOW TO AND