For each pair, which state is more stable?

Pencil on desk vs. raised in the air

Skier at top of mountain vs. bottom

Toaster at 25oC vs. hot toaster

Liquid water at 25oC vs. ice at 25oC

Salt water vs. pure water in contact with solid salt

MgO vs. Mg metal in contact with O2

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

THERMODYNAMICS THERMODYNAMICS Review of Energy and Enthalpy Changes (Ch. 5)

Energy Changes: Heat and WorkEnergy Changes: Heat and Work

Heat = q = energy transferred due to a difference in temperature.

+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.

STATE FUNCTIONSSTATE FUNCTIONS

A state function is a quantity that only does not depend on the process by which the system was prepared

Example: Your altitude (height above sea level) does not depend 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

reactants

products∆E1

2

Energy ChangesEnergy Changes

∆E = Efinal state - Einitial state

∆E (kJ/mol)

O2 (g) 2 O atoms +498.3

Water ice at 25oC -6.0

Water + salt salty water -3.9

Si (s) + O2 (g) SiO2 (s) -908

Which set of reactants/products has the biggest difference in energy?

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?

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) is heated to room temperature at const. P or at const. V

Energy Changes - Heat and WorkEnergy Changes - Heat and Work

∆∆

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 melting

Water boiling

Wood burning

H is a state function – changes are path-independent

H = E + PV (sums and products of state functions are also state functions)

Standard Enthalpy of FormationStandard Enthalpy of Formation

∆Hof:

∆H for making a compound from elements in their standard 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 = ∆Ho

f (products) – ∆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

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 the forward direction under a given set of conditions.

Note: spontaneity has nothing to do with the rate at which a reaction 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?

Useful work can be extracted from a spontaneous process

2 H2 + O2 2 H2O can 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) H > 0

Ba(OH)2•8H2O(s) + 2NH4SCN(s)

Ba(SCN)2(aq) + 2NH3(aq) + 10H2O(l) H > 0

NH4Cl(s) + H2O(l) NH4Cl(aq) H > 0

Spontaneity and Spontaneity and HH

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 chemistry and physics?

ENTROPYENTROPY

A thermodynamic parameter that measures the disorder or randomness in a system.

The more disordered a system, the greater its entropy.

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

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

S = S final - S initial

Which "reactions" have Which "reactions" have S > 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)

ENTROPYENTROPY

Gases have a lot more entropy than solids or liquids.

Reactions that form gases usually have S > 0

S (+ 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)

Adding heat increases entropy

Entropy and heat: S = qrev/T

e.g., for melting ice, S = Hfus/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

Svap = qrev = Hvap

T T

During phase changes temperature and pressure are constant. (Heat transfer is reversible, so H = q = qrev)

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

2nd Law of Thermodynamics:

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

Suniverse = Ssystem + Ssurroundings > 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 only on 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)

So (rxn) = So (products) - So (reactants)

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

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

CRITERIA FOR SPONTANEITYCRITERIA FOR SPONTANEITY

Hsystem < 0Exothermic reactions are usually spontaneous.

SReactions are always spontaneous if

Ssystem + Ssurroundings > 0 (2nd Law)

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

G = H - TS (T = absolute temperature (in K))

Gsystem predicts rxns at constant T and P:

G < 0 SpontaneousG > 0 Not spontaneousG = 0 Reaction at equilibrium

EFFECT OF TEMPERATURE ON SPONTANEITY

G = H - TS

H and S do not change much with temperature, but ∆G does.

H S G Spontaneous?

- +

- - + + + -

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

Hfo 1207.1 635.3 393.5

(kJ/mol)

Calculate the boiling point of bromine

Hovap = 31.0 (kJ/mol)

Sovap = 92.9 (J/mol-K)

STANDARD FREE ENERGY OF FORMATION

Gf = Free energy change in forming one mole of a compound from its elements, each in their standard states.

Standard States:Solid Pure solidLiquid Pure liquid Gas P = 1 atmSolution 1M solutionElements Gf = 0Temperature Usually 25C

Grxn = Gf(prods) - Gf(reactants) Units:kJ/mole.

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 carry out a reaction?

EXTENT OF REACTIONSEXTENT OF REACTIONS

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

What happens under other conditions?

G = G + RTlnQ = G + 2.303RTlog10Q

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

Example: 2NO2 (g) N2O4 (g) Go298 = 5.4 kJ/mol

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

RELATIONSHIP BETWEEN RELATIONSHIP BETWEEN G AND G AND GG

How do concentrations affect Q and G?

QG ( or )Add more reactants

Take away products

Take away reactants

ANALOGY BETWEEN POTENTIAL ENERGY AND FREE ENERGY

At equilibrium point, G = 0 for interconverting reactants products

Note: This does NOT mean Go = 0

slope = 0

Go

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 BETWEENGGAND KAND Keqeq

At equilibrium, G = 0:G = 0 = G + 2.303RTlogQ

= G + 2.303RTlogKeq

G = -2.303RTlogKeq

G > 0 Keq < 1G < 0 Keq > 1G = 0 Keq = 1

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)

Gfo Ag+ 77.1 kJ/mol

Gfo Br 104 kJ/mol

Gfo 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) Equilibrium

Heat in and out is reversible q = qrev

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

RELATIONSHIP BETWEEN RELATIONSHIP BETWEEN G AND WORKG AND WORK

For a spontaneous process,

G = Wmax = The maximum work that can be obtained from a process at constant T and P.

For a non-spontaneous process,

G = Wmin = The minimum work that must be done to make a process go at constant T and P.

Calculating W from Calculating W from GG

What is the maximum work available from the oxidation of 1 mole of octane under standard conditions (P = 1 atm)?

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

Gfo 17.3 0 394.4 237.1

kJ/mol