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For each pair, which state is more stable?Pencil on desk vs. raised in the airSkier at top of mountain vs. bottomToaster at 25oC vs. hot toasterLiquid water at 25oC vs. ice at 25oCSalt water vs. pure water in contact with solid saltMgO vs. Mg metal in contact with O2

In each case, how was energy transferred in going to themore stable state?

THERMODYNAMICSTHERMODYNAMICSReview of Energy and Enthalpy Changes (Ch. 5)

Energy Changes: Heat and WorkEnergy Changes: Heat and WorkHeat = q = energy transferred due to a difference intemperature.+q means heat is added to the system

Work = w = action of force through a distance (often P∆V)+w means work is done on the system

Find one pair in which energy was transferred as heat.Find one pair in which energy was transferred as work.What force was involved?

Find a second pair in which a different force was involved.

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STATE FUNCTIONSSTATE FUNCTIONSA state function is a quantity that only does not dependon the process by which the system was preparedExample: Your altitude (height above sea level) does notdepend on the route you took to class this morning.

State functions are written as uppercase letters (E, H, P,V, T, S…)Changes in state functions are path-independent:

q and w are not state functions but ∆E (= q + w) is a state function

reactantsproducts

∆E12

Energy ChangesEnergy Changes∆E = Efinal state - Einitial state

∆E (kJ/mol)O2 (g) → 2 O atoms +498.3Water → ice at 25oC -6.0Water + salt → salty water -3.9Si (s) + O2 (g) → SiO2 (s) -908

Which set of reactants/products has the biggest difference inenergy?

Which set has the lowest absolute energy?

If we assign O atoms an energy of zero, what is the energy of O2?

Is the lowest energy state always most stable?

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First Law of ThermodynamicsFirst Law of Thermodynamics∆Euniverse = 0

Energy is conserved∆Esystem + ∆Esurroundings = ∆Euniverse

so ∆Esystem = - ∆Esurroundings

Heat and work: ∆Esystem = q + w

and for PV work at const. pressure,

∆Esystem = q – P∆V

w = P!V

+q +q

!V = 0

What are the signs of q and w for each?Does Esystem increase or decrease? How do you know?Which system has higher E at the end?

Dry ice (CO2) isheated to roomtemperature atconst. P or at const. V

Energy Changes - Heat and WorkEnergy Changes - Heat and Work

∆∆

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When natural gas burns:CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(l)

Reaction produces heat and light

Is energy conserved?What are the signs of q and w for the overall reaction?Where does the energy come from?Is energy stored or released when bonds are broken?

Enthalpy (H)Enthalpy (H)∆H = heat transferred at const. P

If (+), endothermic (need to add heat)If (-), exothermic (heat is given off)Classify as endo- or exothermic:

Ice meltingWater boilingWood burning

H is a state function – changes are path-independentH = E + PV (sums and products of state functions are

also state functions)

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Standard Enthalpy of FormationStandard Enthalpy of Formation∆Ho

f:∆H for making a compound from elements in theirstandard states

Standard state is the most stable form (pure solid,pure liquid, or gas at P = 1 atm)For solutes in solution, standard state is usually 1 M

There are tables of ∆Hof

∆Horxn = Σ ∆Ho

f (products) – Σ ∆Hof (reactants)

Standard Enthalpy of Formation of OxidesStandard Enthalpy of Formation of Oxides Reaction ∆Ho

f (kJ/mol)

2 Li + 1/2 O2 → Li2O -598

2 Na + 1/2 O2 → Na2O -414

Ca + 1/2 O2 → CaO -635

Mg + 1/2 O2 → MgO -602

2 Al + 3/2 O2 → Al2O3 -1676

2 Fe + 3/2 O2 → Fe2O3 -824

2 Cr + 3/2 O2 → Cr2O3 -1128

2 Ag + 1/2 O2 → Ag2O -30

2 Cu + 1/2 O2 → Cu2O -167

C + O2 → CO2 -394

1/2 N2 + O2 → NO2 +34

Cl2 + 1/2 O2 →Cl2O +80

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What are the least and most stable oxides in the table?

What periodic trends do you see for the ∆Hof’s of the

oxides?

How do oxidation states relate to ∆Hof?

Predict ∆Hof for Au2O, K2O, and SrO.

Why does Mg burn in dry ice?

SPONTANEOUS REACTIONS

A spontaneous reaction is one that can proceed in theforward direction under a given set of conditions.

Note: spontaneity has nothing to do with the rate at which areaction occurs.

CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g) Spontaneous

CO2(g) → C(s) + O2(g) Not Spontaneous

2 Fe2O3(s) → 4 Fe(s) + 3 O2(g) Not Spontaneous

What determines whether these reactions are spontaneous?

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Useful work can be extracted from a spontaneous process

2 H2 + O2 → 2 H2Ocan be used to drive a rocket

Water in a tower → Water on the ground

Can do work(drive a turbine)

Work must be done to drive a non-spontaneous process2 H2O → 2 H2 + O2 Energy (electrolysis work) must be put in.

Water on the ground → Water in a tower Work is done to pump water

SPONTANEITY AND WORK

Not all spontaneous reactions are exothermic

Examples:At +10°C H2O(s) → H2O(l) Δ H > 0

Ba(OH)2•8H2O(s) + 2NH4SCN(s) →Ba(SCN)2(aq) + 2NH3(aq) + 10H2O(l) Δ H > 0

NH4Cl(s) + H2O(l) → NH4Cl(aq) Δ H > 0

Spontaneity and Spontaneity and ΔΔHH

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QUESTIONSQUESTIONS

Why are some endothermic processes spontaneous?Evaporation of waterDissolution of NH4

+NO3-

What makes an ideal gas expand into a vacuum(q = 0, w = 0)?

What distinguishes "past" from "future" in chemistryand physics?

ENTROPYENTROPY

A thermodynamic parameter that measures thedisorder or randomness in a system.

The more disordered a system, the greater itsentropy.

Entropy is a state function -- its value depends onlyupon the state of the system (not how it got there).

We are usually concerned with the change in entropy(ΔS ) during a process such as a chemical reaction.

ΔS = S final - S initial

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Which "reactions" have Which "reactions" have ΔΔS > 0?S > 0?

Cards arranged in order → Random order after shuffling the deck

Messy dorm room → Cleaned up room

CO2 + H2O + Minerals → Tree

1 mole of gas → 1 mole of gasin a 1 L flask in a 2 L flask

NH4Cl(s) → NH4+(aq) + Cl-(aq)

ENTROPYENTROPYGases have a lot more entropy than solids or liquids.

Reactions that form gases usually have ΔS > 0

ΔS (+ or - ?)

H2O (l, 25oC) → H2O(g)CaCO3 (s) → CaO(s) + CO2(g)

N2(g) + 3 H2(g) → 2 NH3(g)

N2(g) + O2(g) → 2 NO(g)

Au(s) at 298K → Au(s) at 1000K

Ag+(aq) + Cl-(aq) → AgCl(s)

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Adding heat increases entropy

Entropy and heat: ΔS = qrev/T

e.g., for melting ice, ΔS = ΔHfus/273 K (endothermic) ↑ ↑ ↑ + + m.p.= 0oC

Entropy units: J/mol-K (same units as specific heat)

S(1 mole HCl(g)) >> S(1 mole NaCl(s)) why?

S(2 moles HCl(g)) = 2 S(1 mole HCl(g)) “

S(1 mole HCl(g)) > S(1 mole Ar(g)) “

S(1 mole N2(g) at 300K) > S(1 mole N2(g) at 200K) “

ENTROPY and PHASE CHANGESENTROPY and PHASE CHANGES

Sgas >> Sliq > Ssolid

ΔSvap = qrev = ΔHvap T T

During phase changes temperature and pressure areconstant. (Heat transfer is reversible, so ΔH = q = qrev)

Calculate the entropy change when 1 mole of liquidwater evaporates at 100oC (ΔHvap = +44 kJ/mol)

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2nd Law of Thermodynamics:

The total entropy in the universe is increasing.ΔSuniverse > 0

ΔSuniverse = ΔSsystem + ΔSsurroundings > 0

3rd Law:

The entropy of every pure substance at 0K(absolute zero temperature) is zero.S = 0 at 0 K

3rd LAW ENTROPY3rd LAW ENTROPY

Entropy is a state function - its value depends onlyon the initial and final states.

And S = 0 at T = 0 K (3rd Law)

This means we can measure absolute entropy S(not just ∆S)

ΔSo (rxn) = ΣSo (products) - ΣSo (reactants)

N2(g) + 3 H2(g) → 2 NH3(g)So (25oC) = 191.5 130.58 192.5 J/mol-K

ΔSo (rxn) = (2x192.5) - (191.5 + 3x130.58) = -198.3 J/mol-K

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CRITERIA FOR SPONTANEITYCRITERIA FOR SPONTANEITY

ΔHsystem < 0Exothermic reactions are usually spontaneous.

ΔSReactions are always spontaneous if

ΔSsystem + ΔSsurroundings > 0 (2nd Law)

We need a new state function (G) that can predict spontaneityfrom just the system. This is called the Free Energy:

ΔG = ΔH - TΔS (T = absolute temperature (in K))

ΔGsystem predicts rxns at constant T and P:ΔG < 0 SpontaneousΔG > 0 Not spontaneousΔG = 0 Reaction at equilibrium

EFFECT OF TEMPERATURE ON SPONTANEITY

ΔG = ΔH - TΔS

ΔH and ΔS do not change much with temperature,but ∆G does.

ΔH ΔS ΔG Spontaneous?

- +

- -

+ +

+ -

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Is this reaction spontaneous at room temp? At 1100 oC?

CaCO3(s) → CaO(s) + CO2(g) So 92.88 39.75 213.6J/mol K

Δ Hfo −1207.1 −635.3 −393.5

(kJ/mol)

Calculate the boiling point of bromine

ΔHovap = 31.0 (kJ/mol)

ΔSovap = 92.9 (J/mol-K)

STANDARD FREE ENERGY OF FORMATION

ΔGf° = Free energy change in forming one mole of acompound from its elements, each in their standard states.

Standard States:Solid Pure solidLiquid Pure liquidGas P = 1 atmSolution 1M solutionElements ΔGf° = 0Temperature Usually 25°C

ΔG°rxn = ΣΔGf°(prods) - ΣΔGf°(reactants) Units: kJ/mole.

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COMBUSTION OF SUCROSECOMBUSTION OF SUCROSE

C12H22O11(s) + 8 KClO3(s) → 12 CO2(g) + 11 H2O(l) + 8 KCl(s)

Is this reaction spontaneous?

∆Gof(sucrose) = -1544.3 kJ/mol

∆Gof(KClO3, s) = -289.9 ”

∆Gof(CO2, g) = -394.4 "

∆Gof(H2O, l) = -237.1 "

∆Gof(KCl, s) = -408.3 "

How does H2SO4 affect the reaction rate?

What are two possible reasons for using high T to carryout a reaction?

EXTENT OF REACTIONSEXTENT OF REACTIONS

So far we have considered only standard conditionsand ∆Ho, ∆So, ∆Go.

What happens under other conditions?

ΔG = ΔG° + RTlnQ = ΔG° + 2.303RTlog10Q

aA + bB ↔ cC + dD Q = ([C]c[D]d) = REACTION ([A]a[B]b) QUOTIENT

Example: 2NO2 (g) ↔ N2O4 (g) ΔGo298 = −5.4 kJ/mol

Partial pressure of NO2 is 0.25 atm and partial pressureof N2O4 is 0.6 atm. What is ∆G?

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RELATIONSHIP BETWEENRELATIONSHIP BETWEEN ΔΔG AND G AND ΔΔGG°°

How do concentrations affect Q and ΔG?

Q ΔG (↑ or ↓)Add more reactants

Take away products

Take away reactants

ANALOGY BETWEEN POTENTIAL ENERGYAND FREE ENERGY

At equilibrium point, ΔG = 0 for interconvertingreactants ↔ products

Note: This does NOT mean ΔGo = 0

slope = 0

ΔGo

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Nitrogen FixationNitrogen Fixation

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

∆Gof: 0 0 -16.7 kJ/mol

∆Gorxn =

What is the sign of ∆So? ∆Ho?

Would higher or lower P favor products?

" " " " T " " ?

RELATIONSHIP BETWEENRELATIONSHIP BETWEENΔΔGG°° AND KAND Keqeq

At equilibrium, ΔG = 0:ΔG = 0 = ΔG° + 2.303RTlogQ

= ΔG° + 2.303RTlogKeq

ΔG° = -2.303RTlogKeq

ΔG° > 0 Keq < 1ΔG° < 0 Keq > 1ΔG° = 0 Keq = 1

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Solubility EquilibriumSolubility EquilibriumGuess the solubilites of these salts in water:

KClO3(s) → K+(aq) + ClO3-(aq) ∆Go > 0

NaF(s) → Na+(aq) + F-(aq) ∆Go ≈ 0

NaCl(s) → Na+(aq) + Cl-(aq) ∆Go < 0

What is the solubility product Ksp of AgBr?AgBr(s) ↔Ag+(aq) + Br−(aq)

ΔGfo Ag+ 77.1 kJ/mol

ΔGfo Br− −104 kJ/mol

ΔGfo AgBr −96.9 kJ/mol

Spontaneity and EquilibriumSpontaneity and EquilibriumAt −10°C H2O(l) → H2O(s) SpontaneousCan get the system to do work!

At + 10°C H2O(l) → H2O(s) Non- SpontaneousMust do work to get this to go.

But at 0°C H2O(l) H2O(s) EquilibriumHeat in and out is reversible q = qrev

When a chemical system is at equilibrium, reactants andproducts can interconvert reversibly.

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RELATIONSHIP BETWEENRELATIONSHIP BETWEENΔΔG AND WORKG AND WORK

For a spontaneous process,

ΔG = Wmax = The maximum work that can beobtained

from a process at constant T and P.

For a non-spontaneous process,

ΔG = Wmin = The minimum work that must be done tomake a process go at constant T and P.

Calculating W from Calculating W from ΔΔGG

What is the maximum work available from theoxidation of 1 mole of octane under standardconditions (P = 1 atm)?

C8H18(l) + 12.5 O2(g)→ 8CO2(g) + 9H2O(l)ΔGf

o 17.3 0 −394.4 −237.1kJ/mol