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. ∆E universe = ∆E system + ∆E surroundings = 0 ∆E system = -∆E surroundings

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.

E 2 = E 1 + q + w ∆E = E 2 - E 1 = 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(V 2 – V 1 ) ∆E = q + ∆(PV) ∆E = q p + P∆V At constant pressure

Standard Heats of Formation Tables ∆H f 25 o 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 q p = ∆E - P∆V

H 2 O (l)  H 2 O (s) ∆H = Kj/mole T< o C Spontaneous & exothermic H 2 O (s)  H 2 O (l) ∆H = Kj/mole T> o C Spontaneous & endothermic H 2 O (l)  H 2 O (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

S more microstates > S fewer 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

S more microstates > S fewer microstates ∆S system = S final - S initial ∆S system = q reversible T Reversible - means a process that occurs slowly enough for equilibrium to be maintained continuously.

Standard Entropy values Tables S 25 o 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. ∆S universe = ∆S system + ∆S surroundings > 0

Third Law of Thermodynamics- A perfect crystal has zero entropy at a temperature of absolute zero. S system = 0 K

Predicting Relative S values Temperature Changes 273 K295 K298 K S = 31.0 = 32.9 = 33.2 S o increases for a substance as it is heated.

Phase Changes Na H 2 O C (graphite) S initial 51.4 (S) 69.9 (l) 5.7 (s) S final (l) (g) 158 (g) S o increases for a substance as it changes from solid to liquid to gas

Dissolving a Solid or Liquid NaCl AlCl 3 CH 3 OH S o 72.1 (s) 167 (s) 127 (l) S o (aq) 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 O 2 S o (g) = S o (aq) = 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 S o Size (nm) Mass j/(mole x K) Li Na K R Cs

Atomic S o Mass j/(mole x K) HF HCl HBr HI

Allotropes S is greater for the allotrope form that allows the atoms more freedom of motion S o (graphite) = dimensional lattice S o (diamond) = dimensional lattice S o (O 2 gas) = 205 S o (O 3 gas) = ozone

Chemical Complexity Entropy increases with chemical complexity and with the number of atoms in the molecule. NaCl AlCl 3 P 4 O 10 NO NO 2 N 2 O 4 S (g) 240 (g) 304 (g) cyclo CH 4(g) C 2 H 6(g) C 3 H 8(g) C 4 H 10(g) C 5 H 10(g) C 5 H 10(g) C 2 H 5 OH (l) S

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. H 2(g) + I 2(s)  2HI (g) ΔS o Rx = S o P - S o R > 0 1 mole gas to 2 moles gas N 2(g) + 3H 2(g)  2NH 3 (g) ΔS o Rx = S o P - S o R < 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 N 2(g) + 3H 2(g)  2NH 3(g) ΔS o Rx = ΣnS o Products - ΣnS o Reactants = (2 moles NH 3 )(193 J/mole x K) - (1 mole N 2 )(191.5 J/mole x K) - (3 moles H 2 )(130.6 j/mole x K) = J/K