22 September, 1997Chem 1A03E/1E03E THERMOCHEMISTRY (Ch. 6)1 Thermochemistry - Energy of Chemical...

22 September, 1997 Chem 1A03E/1E03E

THERMOCHEMISTRY (Ch. 6)1

ThermochemistryThermochemistry- Energy of Chemical Reactions

ContentsContents:• heat, work, forms of energy• specific heat and energies of phase changes • enthalpy changes in chemical reactions• standard enthalpies of formation • Hess’s law • estimating enthalpies of reaction from Bond Energies



CHEMICAL ENERGYCHEMICAL ENERGY

Chemical bonds are a source of energy• BOND BREAKING - requires energy• BOND MAKING - releases energy

In a chemical reaction:• if more energy is released in forming bonds than is used in breaking bonds then . . . reaction is EXOTHERMICEXOTHERMIC

• if more energy is used in breaking bonds than is released in

forming bonds then . . . reaction is ENDOTHERMIC

Energy is released as HEAT, LIGHT, SOUND, WORK

Energy can be provided by - LIGHT - photochemistry- WORK - electrochemistry- COOLING of surroundings



Energy and ChemistryEnergy and ChemistryENERGYENERGY is the capacity to do is the capacity to do workwork or transfer or transfer heatheat..

HEATHEAT is the form of energy that flows between 2 is the form of energy that flows between 2 samples because of a difference in temperature.samples because of a difference in temperature.

WORKWORK is the form of energy that results in a is the form of energy that results in a macroscopic displacement of matter such as gas macroscopic displacement of matter such as gas expansion or motion of an object expansion or motion of an object (force x distance)(force x distance)

Other forms of energy —Other forms of energy —• light light • electricalelectrical• kinetickinetic

• ChemicalChemical

• gravitational potentialgravitational potential• electrostatic potentialelectrostatic potential



Specific Heat CapacitySpecific Heat CapacitySpecific Heat CapacitySpecific Heat CapacityThermochemistry is the science of heat (energy) flow.

A difference in temperature leads to energy transfer.

The heat “lost” or “gained” is related to a) sample massb) change in T, andc) specific heat capacity by

Specific heat capacity =

heat lost or gained by substance (J)

(mass, g) (T change, K)



Specific Heat CapacitySpecific Heat CapacitySpecific Heat CapacitySpecific Heat Capacity

SubstanceSubstance Spec. Heat (J/g•K)Spec. Heat (J/g•K)

HH22OO 4.1844.184

AlAl 0.9020.902

glassglass 0.840.84

AluminumAluminum

WaterWater



Specific Heat Capacity - an exampleSpecific Heat Capacity - an exampleSpecific Heat Capacity - an exampleSpecific Heat Capacity - an example

If 25.0 g of Al cool from 310 oC to 37 oC, how many joules of heat energy are lost by the Al?

where T = Tfinal - Tinitial = 37 - 310 = -273 Kq = (0.902 J/g•K)(25.0 g)(-273 K)q = -6160 J

negative sign of q heat is “lost by” or transferred from Alnegative sign of q heat is “lost by” or transferred from Al

Specific heat capacity =

heat lost or gained by substance (J)

(mass, g)(T change, K)

0.902 J/g.K

=

heat gain/lost = q = (specific heat)(mass)(T)heat gain/lost = q = (specific heat)(mass)(T)



Heat Transfer and Heat Transfer and Changes of StateChanges of State

Heat Transfer and Heat Transfer and Changes of StateChanges of State

Changes of state involve energyChanges of state involve energy

Ice Water333 J/g333 J/g

(Heat of FusionHeat of Fusion)

Water Vapor 2260 J/g2260 J/g

(Heat of vaporization)(Heat of vaporization)



Heating/Cooling Curve for WaterHeating/Cooling Curve for Water

11 22

33 44

Heat waterHeat water

Evaporate waterEvaporate water

Melt iceMelt ice



CHEMICALCHEMICAL REACTIVITYREACTIVITYCHEMICALCHEMICAL REACTIVITYREACTIVITY

• What drives chemical reactions? How do they occur?

The first is answered by THERMODYNAMICS and the second by KINETICS.

• In Ch. 4 we saw a number of “driving forces” for reactions that are PRODUCT-FAVORED.

• formation of a precipitate

• gas formation

• H2O formation (acid-base reaction)



• Energy transfer also allows us to predict reactivity.

• In general, reactions that transfer energy to their surroundings are “product-favored”.

• How do we describe heat transfer in chemical processes ?

CHEMICALCHEMICAL REACTIVITYREACTIVITYCHEMICALCHEMICAL REACTIVITYREACTIVITY



Heat Energy Transfer in Heat Energy Transfer in Physical & Chemical ProcessesPhysical & Chemical Processes

• COCO2 2 (s, -78 (s, -78 ooC) ---> COC) ---> CO2 2 (g, -78 (g, -78 ooC)C)

Heat flows into the Heat flows into the SYSTEMSYSTEM (solid CO (solid CO22) from the ) from the

SURROUNDINGSSURROUNDINGS in an E in an ENDOTHERMICNDOTHERMIC process.process.

heatheat

SurroundingsSurroundings

SystemSystem



• ENERGY is the capacity to do work or transfer heat.• HEAT is the form of energy that flows between 2 samples because of a difference in temperature.• WORK is the form of energy that results in a macroscopic displacement of matter such as gas expansion or motion of an object (force x distance)

In CO2 sublimation & expansion, the same amount the same amount

of of ENERGY flows from flows from surroundingssurroundings to systemsystem

If expanding gas is enclosed, part of the energy transfer appears in the form of WORK OF EXPANSION

wexp = - PV (for an ideal gas)

If expanding gas is not enclosed, the energy transfer appearsonly as HEATHEAT (CO2 gas gets warm).



FIRST LAW OF THERMODYNAMICSFIRST LAW OF THERMODYNAMICS

q = q = E - wE - w

heat energy transferred

Energy change

work doneby the system

Energy is conserved!Energy is conserved!

OR E = q + wE = q + w

NB - q and w positive when they are

transferred FROM surroundings

TO system


Heat

qsys > 0

System

Workwsys > 0



ENTHALPYENTHALPYENTHALPYENTHALPYMost chemical reactions occur at constant P, so

Heat transferred at constant P is called qp with

qp = H = E - w = E + P V = E+PV)

where H = enthalpy H is defined as E + PV)

H = heat transferred at constant PH = change in heat content of the system

H = HH = Hfinalfinal - H - Hinitialinitial



H = HH = Hfinalfinal - H - HinitialinitialH = HH = Hfinalfinal - H - Hinitialinitial

If Hfinal > Hinitial then H is positive

Process is ENDOTHERMIC

If Hfinal > Hinitial then H is positive

Process is ENDOTHERMIC

If Hfinal < Hinitial then H is negative

Process is EXOTHERMIC

If Hfinal < Hinitial then H is negative

Process is EXOTHERMIC

ENTHALPYENTHALPYENTHALPYENTHALPY



Endo- and ExothermicEndo- and ExothermicEndo- and ExothermicEndo- and Exothermic


Heatqsys > 0

System

ENDOTHERMICENDOTHERMICENDOTHERMICENDOTHERMIC

HeatHeat qqsyssys < 0 < 0


System

EXOTHERMICEXOTHERMIC



But the reverse reaction, the decomposition of water :

H2O(g) + 242 kJ ---> H2(g) + 1/2 O2(g)Endothermic reaction — heat is a “reactant”, H = +242 kJ. This does not occur spontaneously.

Consider the combustion of H2 to form water . .

H2(g) + 1/2 O2(g) ---> H2O(g) 242 kJExothermic reaction — heat is a “product”. H = -242 kJ. This is spontaneous and proceeds readily once initiated.

USING ENTHALPYUSING ENTHALPYUSING ENTHALPYUSING ENTHALPY

BUT . . . Decomposition of water can be made to occur by coupling to another, spontaneous process . . .



How can we How can we make Hmake H22 gas ? gas ?

N. Lewis, N. Lewis,

American Scientist,American Scientist,

Nov. 1995, page 534.Nov. 1995, page 534.

H2O

O2H2

LIGHT

S e m ic o nduc t o r Me t a l

w i r e



Making H2 from liquidliquid H2O involves two steps.

H2O(liq) + 44 kJ H2O(g)

H2O(g) + 242 kJ H2(g) + 1/2 O2(g)---------------------------------------------------H2O(liq) + 286 kJ H2(g) + 1/2 O2(g)

This is an example of HESS’S LAW —

If a reaction is the sum of 2 or more others, the net H is the sum of the H’s of the other rxns.

If a reaction is the sum of 2 or more others, the net H is the sum of the H’s of the other rxns.



Calc. Calc. HHrxnrxn for S(s) + 3/2 O for S(s) + 3/2 O22(g) --> SO(g) --> SO33(g)(g)

knowing thatknowing that

S(s) + OS(s) + O22(g) --> SO(g) --> SO22(g) (g) HH11 = -320.5 kJ = -320.5 kJ

SOSO22(g) + 1/2 O(g) + 1/2 O22(g) --> SO(g) --> SO33(g) (g) HH22 = -75.2 kJ = -75.2 kJ

Hess’s Law - a second example :Hess’s Law - a second example :

The two rxns. add to give the desired rxn.,

S(s) + 3/2 OS(s) + 3/2 O22(g) --> SO(g) --> SO33(g)(g)

so Hrxn = H1 + H2 = -395.7 kJ

HH33 = -395.7 kJ = -395.7 kJ



energy

2

S solid

SO3 gas

direct path

+ 3/2 O

H3 = -395.7 kJ SO2 gas

+O2H1 = -320.5 kJ

+ 1/2 O2H2 = -75.2 kJ

HH33 = -395.7 = -395.7 HH(2+3)(2+3) = -320.5 + -75.2 = -395.7 = -320.5 + -75.2 = -395.7

H along one path = H along one path = H along another pathH along another path

HH33 = -395.7 = -395.7 HH(2+3)(2+3) = -320.5 + -75.2 = -395.7 = -320.5 + -75.2 = -395.7

H along one path = H along one path = H along another pathH along another path



• This equation is valid because This equation is valid because H is a H is a STATE FUNCTIONSTATE FUNCTION

• These depend only on the These depend only on the state of state of the systemthe system and not how it got there. and not how it got there.

• Other Other state functionsstate functions include: include:

V, T, P, energy . . V, T, P, energy . .

H along one path =H along one path =

H along another pathH along another path

H along one path =H along one path =

H along another pathH along another path

— and your bank account!

• Unlike V, T, and P, one cannot Unlike V, T, and P, one cannot

measure absolute H. Can only measure measure absolute H. Can only measure H.H.



Standard Enthalpy ValuesStandard Enthalpy ValuesStandard Enthalpy ValuesStandard Enthalpy ValuesMost H values are labeled Ho

• P = 1 atmosphere ( = 760 torr = 101.3 kPa)

• Concentration = 1 mol/L

• T = usually 25 oC

• with all species in standard states

e.g., C = graphite and O2 = gas

o means measured under standard conditions



- the enthalpy change when 1 mol of compound is formed from elements under standard conditions.

Values: Kotz, Table 6.2 and Appendix KValues: Kotz, Table 6.2 and Appendix K

By definition, Hof = 0 for elements in

their standard states.

Hof = standard molar enthalpy of formationHo

f = standard molar enthalpy of formation

H2(g) + 1/2 O2(g) --> H2O(g)

Hof = -241.8 kJ/mol

H2(g) + 1/2 O2(g) --> H2O(g)

Hof = -241.8 kJ/mol



Using Standard Enthalpy ValuesUsing Standard Enthalpy ValuesUsing Standard Enthalpy ValuesUsing Standard Enthalpy Values

In general, when ALL

enthalpies of formation are known,

Horxn =

Hof (products)

- Hof (reactants)

Calculate Calculate H H of reaction?of reaction?



Horxn = Ho

f (prod) - Hof (react)Ho

rxn = Hof (prod) - Ho

f (react)

Example: Calculate the heat of combustion of ethanol, i.e., Ho

rxn for

C2H5OH(g) + 7/2 O2(g) 2 CO2(g) + 3 H2O(g)

Horxn = { 2 Ho

f (CO2) + 3 Hof (H2O) }

- {7/2 Hof (O2) + Ho

f (C2H5OH)} = { 2 (-393.5 kJ) + 3 (-241.8 kJ) } - {7/2 (0 kJ) + (-235.1 kJ)}

Horxn = -1035.5 kJ per mol of ethanol



• Given by D - the bond dissociation energy

D = energy required to break a bond in a gas phase molecule under standard conditions

e.g. CH4 (g) C (g) + 4 H (g) Hrxn = -1664 kJ = 4 * D(C-H)D(C-H) = 416 kJ per mole of C-H bonds

• D (C-H) (kJ/mol) varies slightly among compounds :

CH4 416 C2H6 392 C3H8 380

C2H4 432 C2H2 445 C6H6 448

Bond Energies Bond Energies (Kotz, sect. 9.4, pp 418-422)(Kotz, sect. 9.4, pp 418-422)

H

H

H

H

• D can be derived from Hrxn for atomization . . .



The GREATER the number of bonds (bond order) the HIGHER the bond dissociation energy

The GREATER the number of bonds (bond order) the HIGHER the bond dissociation energy

BOND D (kJ/mol) (Bond Energy) H—H 436 C—C 347 C=C 611 CC 837

N—N 159 NN 946see table 9.5 for Dissociation Energies of other bonds.

• D is similar for same bond in different molecules• Average values over many compounds are tabulated• Bond energy depends on bond order

Bond Energies Bond Energies



Using Bond EnergiesUsing Bond Energies

• Estimate the energy of the reaction

H—H + Cl—Cl ----> 2 H—ClH—H + Cl—Cl ----> 2 H—ClH—H + Cl—Cl ----> 2 H—ClH—H + Cl—Cl ----> 2 H—Cl

Net energy = Hrxn = energy required to break bonds - energy evolved when bonds are made



H—H = 436 kJ/molH—H = 436 kJ/molCl—Cl = 243 kJ/molCl—Cl = 243 kJ/mol H—Cl = 431 kJ/molH—Cl = 431 kJ/mol

H—H = 436 kJ/molH—H = 436 kJ/molCl—Cl = 243 kJ/molCl—Cl = 243 kJ/mol H—Cl = 431 kJ/molH—Cl = 431 kJ/mol

• Sum of H-H + Cl-Cl bond energies = 436 kJ + 243 kJ = +679 kJ

• 2 mol H-Cl bond energies = 862 kJ

• Net = H = +679 kJ - 862 kJ = -183 kJ

THEREFORE, , Hf for H-Cl is ? ? ?? ? ?

Estimating Hrxn for H—H + Cl—Cl 2 H—Cl



• Is the reaction exo- or endothermic?

Energy for bond breaking:

4 mol O—H bonds = 4 (464 kJ)

2 mol O—O bonds = 2 (138 kJ)

TOTAL = 2132 kJ

Energy from bond making :

1 mol O=O bonds = 498 kJ

4 mol O—H bonds = 4 (464 kJ)

TOTAL = 2354 kJ

EXAMPLE 2: Estimate the energy of the reaction

2 H—O—O—H ----> O=O + 2 H—O—H

• Which is larger: energy req’d to break bonds . . .

or energy evolved on making bonds?



2 H—O—O—H ---->

O=O + 2 H—O—H

More energy is evolved on More energy is evolved on making bonds than is making bonds than is expended in breaking expended in breaking bonds.bonds.

More energy is evolved on More energy is evolved on making bonds than is making bonds than is expended in breaking expended in breaking bonds.bonds.

The reaction is exothermic!The reaction is exothermic!The reaction is exothermic!The reaction is exothermic!

Net energy = +2132 kJ - 2354 kJ = - 222 kJNet energy = +2132 kJ - 2354 kJ = - 222 kJ



Enthalpies of Reaction from Bond EnergiesEnthalpies of Reaction from Bond Energies

REACTANTS

GaseousAtoms

PRODUCTS

ENDOTHERMICENDOTHERMIC

Bond Breaking costs morethan is gained by Bond Making

Bond Making releases more Ethan required for Bond Breaking

PRODUCTS

REACTANTS

EXOTHERMICEXOTHERMIC

GaseousAtoms



Key Concepts from Chapter 6: ThermochemistryKey Concepts from Chapter 6: Thermochemistry

• heat transfer - specific heat• phase transitions - heats of fusion, vaporization, etc

• First law of thermodynamics E = q - w

• endothermic versus exothermic reactions• enthalpy change in chemical reactions• Hess’s law

• standard molar enthalpies of formationHrxn = Hf(products) - Hf (reactants)

• bond energies Hrxn = D(bonds broken) - D(bonds made)

22 September, 1997Chem 1A03E/1E03E THERMOCHEMISTRY (Ch. 6)1 Thermochemistry - Energy of Chemical...

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Transcript of 22 September, 1997Chem 1A03E/1E03E THERMOCHEMISTRY (Ch. 6)1 Thermochemistry - Energy of Chemical...