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THERMOCHEMISTRY

The Nature of Energy

I. Energy is hard to define precisely.

2. The book defines energy as the capacity to do work or produce heat.

3. We will concentrate on the heat transfer that accompanies chemical processes.i-AI.V cF CvJSe:12vl1flV.,.

4. The Or ENon 7 states that energy can be converted from one form to another but itcan be neither created nor destroyed.

is energy due to position or composition.

6. The !<lrJE'T/C CAiEJ2..c, l of an object is energy due to the motion of the object and it depends onthe mass of the object and its velocity. It is calculated using:

K'E:: t..,.,.",-,1

7. Energy can be converted from one form to another form.

8. A lot of the energy transfer in chemistry takes place in the form of heat.

9. T9J'I(Ji;'1!,lTL,'liFsubstance.

is a property that reflects the random motions of the particles in a particular

10. 11m, involves the transfer of energy between two objects due to a temperature difference.

I I. A smrr- &';.JCTIO,j refers to a property of the system that depends only on its presentstate. A state function (property) does not depend in any way on the system's past of future. It doesnot depend on pathway.

12. A change in a state function in going from one state to another state is independent of the particularpathway taken between the two states.

13. Energy is a state function, but work and heat are not.

Chemical Energy

I. Let's look at the combustion of methane:

C 1-1 'f (li) 1 ) O,:J (e) -7 C()ij (g'I-lJllh () 1) -+ b./a:-CJ (lfc:n;)..J J

2. The stiTt!>, is the part of the universe on which we wish to focus attention.

<'I,}~I'iA ';j f)J tJ(,..{3. The..) "-~ Include everything else in the universe.

4. When energy flows out of the system we have an i'xoTH~IJ 1/C process. When it flows into thesystem it is an t:t.f{A:.J'mElZml C. process.

OVER

r •

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5. The difference in potential energies between the reactants and the products causes energy tlow.

6. The energy gained by the surrounding must be equal to the energy lost by the system.

7. In any exothermic reaction, some of the potential energy stored in chemical bonds is being convertedto thermal energy.

8. An energy diagram can be used to help visualize the energy flow.

:lrw-t 0..-{y'\ve GI,/

).w~ UzOI•.•u-i ca~

9. i\ r>E represents the change in potential stored in the bonds of the products as compared with thebonds of the reactants.

,"'IE oj:: .$'1s'ffJ'-1()a:.lt:~rSES

I/'!.PE

J

ctJfJlG'f IS

{<EL£YlSt-o TUmE ~lU.l;.u.);N (, .s

AS H~")'4T

FtiENfI,\LENeZ(,f

10. It's the difference between the energy required to break the bonds in the reactants and the energyreleased when the bonds in the products are formed.

11. The study of energy and its interconversions is called melMOD" "-'I) "1' CS .1=1,2~T U~W cF

12. The law of conservation of energy is often called the THelM ()D'INA M I(Sof the universe is constant.

: the energy

13. The :INfEi.JJ:\l l:-74t~(,'" (E) of a system can be defined most precisely as the sum of the kineticand potential energies of all the "particles" in the system.

14. The internal energy ofa system can be changed by a flow of work, heat, or both:

~E z: H6'\i + Wt'.z..~

15. Thermodynamic quantities always consist of two parts: a 1J(/f'1(3t::'I. ,giving the magnitude of thechange, and a S/l:.;r,J , indicating the direction offlow.

16. The sign reflects the system's point of view.

17. The J{)..:~E(:r) is the fundamental Sf unit for energy:

1.~IVl L

v.-z..,)

..J

P-V Work

I. A common typeof work associated with chemical processes is work done by a gas expansion or workdone to a gas, compression.

2. Pressure is force per unit area:FA

3. For an expanding gas, l\V is positive.ex 1",1"-1':>,,-.; c iXc:.s

.$(;~l,;v.) /I.JGJu,'-lENTHALPY AND CALORIMETRY

Enthalpy

I. Enthalpy, H, is defined as:

:::

2. Since internal energy, pressure and volume are all state functions, enthalpy is also a state function.

3. At constant pressure (where only PV work is allowed), the change in enthalpy, l\H, of the system isequal to the energy flow as heat.

4. Heat of reaction and change in enthalpy are interchangeable terms.

5.

6. If the products have greater enthalpy than the reactants, l\H will be positive. Heat will be absorbed bythe system, the reaction is endothermic.

7. If the enthalpy of the products is less than that ofthe reactants, l\H will be negative. Heat will bereleased by the system; the reaction is exothermic .

••••• The overall reaction in commercial heat packs can be represented as:

l\H = ·1652 kJ

a Determine the amount of heat released when 4.00 mol Fe is reacted with excess O2

b. Determine the amount of heat released when 1.00 mol Fe20] is produced.

OVER

c. Determine the amount of heat released when l.OOg Fe is reacted with excess O2•

d. Determine the amount of heat released when IO.Ogof Fe and 2.00 g O2 are reacted.

--LJ- it-V c re ImJ Fe ..I(,~.~..tJ- 7J. ()Lzr" - s-- ,- 1",~fr;,

.:.SS . J') j-- 1:

(L '0 /(oSl.t.J"c).en ~ j' .,

"Ie,,) '" I .,...~ "" -. .• - J '/. tf -1-J3 J. .COj 0;7 ::> i ()J •..•.•, '(;)

Calorimetry

l. A CI1teeW1E1DZ is the device used experimentally to determine the heat associated with achemical reaction.

2. Calorimetry, the science of measuring heat, is based on observing the temperature change when a bodyabsorbs or discharges energy s heat.

3. The heat capacity ofa substance:

11c.")\T A,.3:.i0V3EDL s:p.j(iutl')E .:tN Tf¥nPt~t.m.Ie[

4. Since the amount of heat required to raise the temperature of a substance by one degree Celsiusdepends on the amount of substance present, the amount must be specified.

5. !,I-:t-Cifl( Ht:AT C~i}ttr"lis the energy required to raise the temperature of one gram of a substance byone degree Celsius. J...... ~

JOe. d JK..6. il,}(,tuL Hen CM4cnl is the energy required to raise the temperature of one mole of substance by

one degree Celsius.

7. Units are:.J .7

tW.("( ~ r-II-\l:t

t.

8. Metals have a much lower heat capacity than water.

9. The measurement ofheat using a simple calorimeter is an example ofUk.Jsl})JT - i12.f:'-~ZE Ct~l Ci2..1t-; CTlt'f since the pressure (atmospheric) remains constant during

the process.

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0.·"c '".

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10. Constant pressure Calorimetry is used in determining the change in enthalpy (heat of reaction) forreactions occurring in solution.

II. The heat released by a reaction an be determined by:~ ~ I/EYjT

S s: ,s-pa::,F Ie HOlT C4P.ic 17'1

(VJ 7 mm OF ~""~Lncj,]

llT -: Tp - ~ .••••• A 46.2g sample ofCu is heated to 9SAoC and then placed in a calorimeter containing 75.0gwater at 19.6°C. The final temperature of the metal and the water is 21.g0C. Calculate the specificheat capacity ofCu, assuming that all the heat lost by the copper is gained by the water.lt~I~£r)~ - g us r:

.sn) ll'T:: - ,st')) AI

eLl S) ( 7oS .0) ( .1/. '8"- /(;. ~ )

(I.If",). ) (.:l/.f - <Jr. If) ~ - S·

J' =••••• Consider the reaction:

2 HCI (aq) + Ba(OHh (aq) ~ BaCh (aq) + 2 H20 (I) AH =-1 18 kJ

Calculate the heat when 100.0 mL of 0.500 M HCI is mixed with 300.0 mL of 0.500 M Ba(OHh.Assuming that the temperature of both solutions was initially 2S.0oC and that the final mixture has a.mass of 400.0g and a specific heat capacity of 4.18 J/gOC, calculate the final temperature of themixture .

• /00 L HeR . .scothvfl-tC(' -, 1/8" ..is""I I i HCQ ".Jm,-t HCt

• 3COL&'iCOH)l Sco..,.~ J:.t\(Oli)L '-II~.-t.T\ I L &.. (O~)l / 1'rI{~ fx..(OH)l

z:

- 17. 71..:r

). CiS A:J ;: ('-1.16') (1co .())( 7f - :JS () 'oC )

.J, {j5 'flU:' J : (V"~-)(VOO, O)(Tj - .)S"(l)

.J . liS of, 10

(~.I~)('-tee ,c)..;- ,]5.() t:

12. Calorimetry can also be performed at N;.J.JTf.1AIT VCLur.}c .

13. To study reactions under this condition, they are studied in a bomb calorimeter.

OVER

. "

••••• A O.1964g sample of quinone (C6H402) is burned in a bomb calori~eter that has a heatcapacityof 1.56 kJi'C. The temperature increases by 3.2oC .. Calculate the energy of combustion of quinone pergram and per mole ..

HEGT cS~lrJl:l) /3'1 CI:tlC.iil"lcn.:ll.:. - He:\T L.C~T 13'( Qv,'""o ••..J(7

(I.~{p . ~~ )( 3. L GC)

J../.l}{1L J ;r

_ ~S.4J:J

5108. ()t/5

Ir>\'--€.J 7tf7

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HESS'S LAW

I. The change in enthalpy in going from some initial state to some final state is independent of the. pathway.

2. In going from a particular set of reactants to a particular set of products, the change in enthalpy is thesame whether the reaction takes place in a single step or in a series of steps.

3. This is known as HESS 1$ LA tt0

4. Let's look at the oxidation of nitrogen to produce nitrogen dioxide:

In multiple steps:

iJ'Ll1) + OJ <;9) -"7 dIJO ~)

.:) AI 0 ~) ~ O,;} (5) ~ :; JO <J. lJ;

/)H ~ I ro .-K.r~H':; - /I;)l:...J-

Atl -

The sum of the two steps gives the net, or overall, reaction.

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t.. -.

Aa·. ", ..

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CHARACTERISTICS OF ENTHALPY CHANGES

I. If a reaction is reversed, the sign of 6H is also reversed.

2. The magnitude of 6H is directly proportional to the quantities of reactants and products in a reaction.Jfthe coefficients in a balanced reaction are multiplied by an integer, the value of 6H is multiplied bythe same integer.

3. l\H is an EnBJSIVf property.

• •••• Combustion reactions involve reacting a substance with oxygen. When compounds containingcarbon and hydrogen are combusted, carbon dioxide and water are the products. Using the enthalpies ofcombustion for C4H4 (-2341 kJ/mol), C6HS (-2755 kJ/mol) and H2 (-286 kJ/mol) calculate the l\H for thereaction:

Calculate the l\H from the following data:

C.•H.•(g) + 5 O2 (g) ~ 4 CO2 (g) + H20 (I)

C4H. (g) + 6 O2 (g) ~ 4 CO2 (g) + 4 H20 (I)

l\H,omb= -2341 kJ/mol

l\H,omb= -2755 kJ/mol

l\Hcomb= -286 kJ/mol

C 'I ~'l lj 7 + 5 O;L (5 j ~ L{ (U CJ) ~ ~J{0 ((1)

(,;{) )./;'{S) .1- (.1) t02~) -"7(';)/-/,0(1)

1 CO~<y) -f I{ Hz O(R) -*? Cvlf fJI -I t OJ. fJ)

Cl{1Iv ~) ~ ol M '1) ~ (I( lit 5)

~H;; - d3l/1 .t.r

l-Z) bll-f= - Sf, l Er:

AI-! = ~7SS)(J

AI-{ :: - /SK JeT

OVER

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••••• The bombardier beetle uses an explosive discharge as a defensive measure. The chemical reactioninvolved is the oxidation of hydroquinone by hydrogen peroxide to produce quinone and water:

Calculate the AH from the following data:

MI = 177.4 kJ/mol

Hz (g) + O2 (g) -+- H20z (aq) AH = -191.2 kllmol

dH = -241.8 kllmol

dH = -43.8 kllmol

c.Wv (o.~L -" CiJ I-Ic, 0iJ. +1-0

I~2- 0 - -7 I-I~ r o·01.01.

;) Hd -+ ()~ -7 () fIz 0

d 111.-0 --71 clilL 0

6 1-/:: I 7 7.</ Er:

() 1-1:- Itl/. J ~J

AH~ - l/fJ.' .t.r:A 1-1 = - f Z {; .Ls:

Ll i-I .:- dOJ.'- ..tT

I'I-

STANDARD ENTHALPIES OF FORMATION

I. Some reactions don't lend themselves well to study within a calorimeter.

2. We can calculate 6H values from standard enthalpies of formation.

3. The ;;TJ;Jof4l1> eJ.ltHi)LP1 t'r" r0 ••r):1n~f~ (I\~t) is the change in enthalpy that accompanies theformation of one mole of a compound from its elements with all substances in their standard state.

4. A degree symbol on a thermodynamic symbol indicates that the corresponding process has been carriedout under standard conditions.

5. The SniJD"JI) ~n-ln= for a substance is a precisely defined reference state,

6. Because thermodynamic functions often depend on the concentrations (or pressures) of the substancesinvolved, we must use a common reference state to properly compare the thermodynamic properties of twosubstances.

7. This is especially important because, for most thermodynamic properties, we can measure onlyC.HArJ(, ~ j in the property.

Conventional Standard States

1. For a compound - the standard state of a gaseous substance is a pressure of exactly one atmosphere.

2. For a compound - the standard state for a pure substance in a condensed state (liquid or solid) is thepure liquid or solid.

3. For a substance present in a solution, the standard state is a concentration of exactly 1 M.

4. The standard state for an element is the form in which the element exists under the conditions of oneatmosphere and 25°C.

Using Heats of Formation

I. Enthalpies of formation are always given per mole of product with the product in its standard state.

2. Enthalpy is a state function so we can invoke Hess's law and choose any convenient pathway fromreactants to products and then sum the enthalpy changes along the chosen pathway.

3. The enthalpy change for a given reaction can be calculated by subtracting the enthalpies of formation ofthe reactants from the enthalpies of formation of the products:

2 1\ f c~I-l~+ fJ,i)vciSDH~"{'It' r-1UL[-'S c;:: e-ILH P';?~'JL'(!

il, I"XilC'> cF (5"1\(+-4 j4:-j.)..:·-r/~A..l·;-4. Elements are not included in the calculation because elements require no change in form.

OVER

" .. - ..••••• Use the values of ~Hfo in Appendix 4 to calculate ~Ho for: .'

'..,;.:

C2HsOH (I) + 3 O2 (g) - 2 CO2 (g) + 3 H20 (g) ,

," .;l,8 -kr ¢ -393,StJ -JlfJ..tI

r La (- 3(;3. S) ,'"3 ( - cl«a. ) J - (-.l- "1S- )

= - I,) 3::> --t_:T

••••• The space shuttle orbiter utilizes the oxidation of methyl hydrazine by dinitrogen tetroxide forpropulsion:

4 N2H3CH) (I) + 5 N204 (I) - 12 H20 (g) + 9 N2 (g) + 4 CO2 (g)s'l ,"2.0 - L'll.. 0 - -3'I~. S-

Calculate ~Ho for the reaction.

-..'~ ,