• First Law of Thermodynamics-The total amount of energy in the universe is constant.

• Second Law of Thermodynamics- All real processes occur spontaneously in the direction that increases the entropy of the universe.

• Third Law of Thermodynamics- A perfect crystal has zero entropy at a temperature of absolute zero.

The Laws of Thermodynamics

First Law of Thermodynamics

The total amount of energy in the

universe is constant.

∆Euniverse = ∆Esystem + ∆Esurroundings = 0 ∆Esystem = -∆Esurroundings

The internal energy ( ∆E ) of a system is the sum of the kinetic and potential energy of all its particles.

A spontaneous change occurs when a chemical reaction proceeds towards equilibrium. Non-spontaneous processes require a continuous input of energy. This does not mean a spontaneous change is instantaneous.

E2 = E1 + q + w

∆E = E2 - E1 = q + w

q = heat transfer (+) heat energy transferred from

surroundings to the system (-) heat energy transferred from system to the surroundings w = work (+) work done on system by surroundings (-) work done on surroundings by system

Work in chemistry is pressure-volume changes

w = - ∆(PV) usually constant pressure It is negative if energy is required to increase

the volume of the system

w = - P∆V = - P(V2 – V1)

∆E = q + ∆(PV) ∆E = qp + P∆V At constant pressure

• Standard Heats of Formation

Tables ∆Hfo @ 25o C Kj/mole

Hess’s Law of heat summation = The enthalpy change for the overall reaction equals the sum of the

enthalpy changes for the individual steps.

Endothermic + ∆H Exothermic - ∆H usually spontaneous

but not always

Enthalpy ∆H Ξ qp = ∆E - P∆V

H2O(l) H2O(s) ∆H = - 6.02 Kj/mole T<oC

Spontaneous & exothermic

H2O(s) H2O(l) ∆H = + 6.02 Kj/mole T>oC

Spontaneous & endothermic

H2O(l) H2O(g) ∆H = 44.0 Kj/mole

Spontaneous & endothermic

Enthalpy is not an absolute predictor of spontaneity

In thermodynamic terms, a change in the freedom of motion of particles in a system and in the dispersal of the energy of motion is a key factor determining the direction of a spontaneous process

Why more freedom of particle motion – energy of motion becomes dispersed (or spread over more quantized energy levels)

Localized has less freedom of motion

Dispersed has more freedom of motion

Microstates• Systems with fewer microstates have

lower entropy• Systems with more microstates have

higher entropy

Phase changes S L G

Dissolution Crystalline solid + liquid water aqueous ions

Chemical Change Crystalline solid gases + aqueous ions

Smore microstates > Sfewer microstates

∆S system = S final - S initial

Entropy is a thermodynamic quantity related to the number of ways the energy of a system can be dispersed through the motion of the particles

Entropy

• Standard Entropy values

Tables So @ 25o C joules/(mole X K)

Hess’s Law of summation = The entropy change for the overall reaction equals the

sum of the entropy changes for the individual steps.

Second Law of Thermodynamics

All real processes occur spontaneously in the direction that increases the entropy of the universe.

∆Suniverse = ∆Ssystem + ∆Ssurroundings > 0

Third Law of Thermodynamics-

A perfect crystal has zero entropy at a temperature of absolute zero.

Ssystem = 0 @ 0 K

Predicting Relative S values

• Temperature Changes

273 K 295 K 298 K

S = 31.0 = 32.9 = 33.2

So increases for a substance as it is heated.

• Phase Changes

Na H2O C (graphite)

Sinitial 51.4 (S) 69.9 (l) 5.7(s)

S final 153.6 (l) 188.7(g) 158 (g)

So increases for a substance as it changes from

solid to liquid to gas

• Dissolving a Solid or Liquid

NaCl AlCl3 CH3OH

So 72.1 (s) 167(s) 127(l)

So(aq) 115.1 -148 132

Ionic solids dissolve in water. Crystals break down increasing freedom of motion dispersed over more microstates.

Hydrated ions, like the Al(aq)+3 ion, make a more

organized unit resulting in a negative entropy change.

Positive ∆S values are very small for a liquid dissolved in another liquid.

• Dissolving a Gas in Water

O2 So(g)

= 205.0

So(aq) = 110.9

When a gas is dissolved in a liquid ∆S is negative.

less freedom

When a gas is dissolved in a gas ∆S is positive.

more freedom

• Atomic Size or Molecular Complexity

(same phase)

Atomic Atomic So Size (nm) Mass j/(mole x K)

Li .205 6.9 29.1

Na .223 23.0 51.4

K .277 39.1 64.7

R .298 85.5 69.5

Cs .334 132.9 85.2

Atomic So Mass j/(mole x K)

HF 20.0 173.7

HCl 36.5 186.8

HBr 80.9 198.6

HI 127.9 206.3

Allotropes

• S is greater for the allotrope form that allows the atoms more freedom of motion

• So (graphite) = 5.96 3 dimensional lattice

• So (diamond) = 2.44 3 dimensional lattice

• So (O2 gas) = 205• So (O3 gas) = 238.8 ozone

Chemical Complexity

Entropy increases with chemical complexity and with the

number of atoms in the molecule.

NaCl AlCl3 P4O10 NO NO2 N2O4

S 72.1 167 229 211(g) 240(g) 304(g)

cyclo

CH4(g) C2H6(g) C3H8(g) C4H10(g) C5H10(g) C5H10(g) C2H5OH(l)

S 186 230 270 310 348 293 161

• Number of moles

If the number of moles of gas increases then ∆S

is usually positive. If the number of moles

decreases then ∆S is usually negative.

H2(g) + I2(s) 2HI(g) ΔSoRx = So

P - SoR > 0

1 mole gas to 2 moles gas

N2(g) + 3H2(g) 2NH3 (g) ΔSoRx = So

P - SoR < 0

4 mole gas to 2 moles gas

Remember you cannot predict the sign of entropy unless the reaction involves a change

in the number of moles of gas

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

ΔSoRx = ΣnSo

Products - ΣnSoReactants

= (2 moles NH3)(193 J/mole x K)

- (1 mole N2)(191.5 J/mole x K)

- (3 moles H2)(130.6 j/mole x K)

= - 197 J/K