Thermodynamics
By Alex Weber and Lee Cheung
Standard Thermodynamic
Conditions

25 C°, and 1 barr = 100kPa ≈ 1 atm
Specific Heat (C)
 Energy
required to raise a gram of material
by 1° C
 C is a constant based on material and state
of the matter
 Units in J/g° C
 C = 4.18 J/g° C (For Water)
Heat Transferred (q)
q= mC(ΔT) where C is specific heat, ΔT is
the change in temperature (in ° C), and m is
mass in grams
 Units for q in Joules
 +q means heat is transferred to the system
 -q means heat is transferred from the system
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Bond Energies
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q = Σ Reactants Bond Energy – Σ Products Bond
Energy (note: reactants – products)
Endothermic reactions need heat/energy added to
progress
Breaking bonds requires heat/energy to progress
Exothermic reactions give off heat/energy during
the reaction
Forming bonds gives off heat/energy
Potential energy in an elemental state = 0
Enthalpy (ΔH)
 ΔH
is the heat transferred into a system per
mole or per gram
 ΔH = q/m= change in potential energy from
products to reactants
 ΔHrxn = ΣHf products - ΣHf reactants
 -ΔH = exothermic
 +ΔH = endothermic
 Units (J/g or J/mol)
Hess’s Law

If a reaction equals the sum of a series of
reactions, then the overall ΔH equals the
sum of ΔH from each individual reaction
 Reverse reactions = reverse signs
 If you change the coefficients of a reaction
by a certain factor, then change ΔH by the
same factor
Entropy (ΔS)
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ΔS is the measure of randomness of molecules
ΔS = the change in S from products to reactants
Gases are the most random and have the highest
entropy, solids the lowest
Product favored reactions have higher entropy
Units of J/K*mol
Calorimetry
 Mmetal
cΔT = mwater cΔT
 Mmetal c(Tf-Ti) = mwater c(Tf-Ti)
 q reaction = - (q water + q bomb)
Changes in State
 q=
mHf or q= mHv (Hf = Heat of fusion,
Hv = Heat of vaporization)
 Heat of fusion = heat required to melt a
substance into liquid
 Heat of vaporization = heat required to
vaporize substance into gas
Gibbs Free Energy
ΔG = ΔH - TΔS (T is in K)
 When ΔG is negative reaction is
spontaneous and vice versa
 Threshold Energy = when ΔG = 0, equation
is at equilibrium
 Spontaneous reactions favor products

Gibbs Free Energy Cont.
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∆G = ∆G˚ + RT lnQ
∆G˚ = -RT lnK (at equilibrium), where R=8.314
J/(mol*K) and T = Temperature (K)
K = Thermodynamic Equilibrium Constant
Q= reaction quotient = K (at equilibrium)
When ∆G˚ < 0 and K > 1 Reaction is product
favored (spontaneous)
When ∆G˚ = 0 and K = 1 Reaction is at
equilibrium
When ∆G˚ > 0 and K < 1 Reaction is reactant
favored (non-spontaneous)
Spontaneity
Considering T= ΔH/ΔS when
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ΔS < 0

ΔS > 0
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
ΔH < 0
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Spontaneous at
Low Temps
Always
Spontaneous
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ΔH > 0
Never
Spontaneous
Spontaneous
at Higher Temps