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• 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
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First Law of Thermodynamics
The total amount of energy in the
universe is constant.
∆Euniverse = ∆Esystem + ∆Esurroundings = 0 ∆Esystem = -∆Esurroundings
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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.
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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
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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
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• 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
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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
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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
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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
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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
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• 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.
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Second Law of Thermodynamics
All real processes occur spontaneously in the direction that increases the entropy of the universe.
∆Suniverse = ∆Ssystem + ∆Ssurroundings > 0
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Third Law of Thermodynamics-
A perfect crystal has zero entropy at a temperature of absolute zero.
Ssystem = 0 @ 0 K
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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.
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• 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
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• 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.
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• 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
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• 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
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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
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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
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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
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• 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
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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