THERMODYNAMICS
Spontaneous Processes
 Chemical thermodynamics is concerned with energy
relationships in chemical reactions
o We consider enthalpy
o We consider disorder (entropy)
 First Law of Thermodynamics : energy is conserved or energy
is neither created or destroyed
o E=q+w
o Where E is the change in internal energy, q is the heat
absorbed by the system, and w is the work done
 Any process that occurs without intervention is spontaneous
o Spontaneous reactions have direction
o Drop an egg, it breaks spontaneously. The reverse
reaction is not spontaneous
 A process that is spontaneous in one direction is not in the
opposite direction
o Direction can depend on temperature
o Water to ice is spontaneous at T0 C
o Ice to water is spontaneous at T 0 C
Reversible and Irreversible Processes
 A reversible process is one that can go back and forth along the same
path
 Chemical systems at equilibrium are reversible
o They interconvert between products and reactants
o Water and ice are at equilibrium at 0C
o There is only one reversible path between 2 states of a system
 An irreversible process is one that can not be reversed to restore the
original state
o To get back to the original state, the species must follow a
different pathway
o In any spontaneous reaction, the path between reactants and
products is irreversible
 Thermodynamics gives us the direction of a process
o It does not predict the rate of the process
o Can an endothermic process be spontaneous?
Entropy and the Second Law
 Spontaneous Expansion of A Gas
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o
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Consider 2 one liter flasks connected by a closed stopcock
One flask is evacuated and the other contains 1 mole of gas
Open the stopcock while maintaining constant temperature
Final state: both flasks contain .5 mole of gas
The expansion is isothermal. No work is done. No heat is
transferred
 Why is this process spontaneous?
 Why is the reverse process not spontaneous?
 Consider 2 molecules when the stopcock is opened. They can
randomly distribute into 1 of 4 possibilities. All in one flask, all in the
other, one in each. The probability that both molecules will be in one
flask is given by (1/2)n where n is the number of molecules. So in this
case ¼.
 For 1 mole of gas n= 6.02 X 1023 so there is essentially no chance that
all molecules will be in one flask
 When the molecules spread out, there was an increase in the
randomness or disorder of the system
 Processes that increase the disorder of the system tend to be
spontaneous
Entropy
 Consider the melting of ice
o In ice, the molecules are held in a rigid lattice
o As the ice melts, the molecules are more randomly distributed
 Consider dissolving a KCl crystal
o Solid KCl has ions in a highly ordered arrangement
o When the crystal dissolves, the ions have more freedom. They
are more random
o However, the water molecules are more ordered. Since, some
molecules are used to hydrate the crystal
o This example involves disordering and ordering
o The disordering predominates
 Entropy, S , is the thermodynamic term that describes the degree of
disorder of a system
o The more disordered or random the system, the larger S
o Entropy is a state function. It is independent of path
o For a system, S = Sfinal - Sinitial
o If S 0 randomness increases; If S 0, order increases
 Suppose a system changes reversibly between state 1 and state 2
o The change in entropy is given by: S = qrev/T
o Where q is the heat added reversibly to the system
o A phase change occurs at constant T with the reversible
addition of heat
The Second Law
 In any spontaneous process, the entropy of the universe must increase
 The change in entropy of the universe is given by the sum of the
change in entropy of the system and the surroundings
 Suniverse = Ssystem + Ssurroundings
 For a reversible process, the sum equals zero
 For a spontaneous process(irreversible), the sum is greater than zero
 Entropy is not conserved. The entropy of the universe is continually
increasing. The second law states that the entropy of the universe must
increase in a spontaneous process
Molecular Interpretation of Entropy
 The entropy of a system indicates it’s state of disorder
o A gas is less ordered than a liquid, which is less ordered than a
solid
o Any process that leads to an increase in gas molecules increases
entropy
 Individual molecules have degrees of freedom associated with motion
within the molecule
 There are 3 modes of atomic motion
o Translational
 Movement from one point to another
o Vibrational
 Shortening and lengthening of bonds,change in bond
angles
o Rotational
 The spinning of a molecule about an axis
 Energy is required for the molecule to translate, vibrate or rotate
o These are forms that a molecule stores energy
o The more stored energy increases entropy
 In a perfect crystal at 0K there is no molecular motion
o This is a state of perfect order
o Entropy is zero
o This is the third law of thermodynamics
 Entropy will increase as we increase the temperature from 0K
o Molecules gain vibrational motion
o Degrees of freedom increase
o Entropy changes dramatically at phase changes
 In general , entropy will increase when:
o Liquids or solutions are formed from solids
o Gases are formed from liquids or solids
o The number of gas molecules increases
o The temperature is increased
Calculation of Entropy Changes
 Standard molar entropy, S, is the molar entropy of a substance in its
standard state
 Units are J/mol K
 Observations about S values
o Standard molar entropies of elements are not zero
o S values are greater for gases than for liquids or solids
o S tend to increase with increasing molar mass
o S tends to increase with the number of atoms in the formula of
a substance
o For a chemical reaction that produces n moles of products and
has m moles of reactants
 S =  nS(products) - mS (reactants)
Gibbs Free Energy
 Reactions with a large negative H tend to be spontaneous
 How can we use S and H to predict whether a reaction is
spontaneous
 The Gibbs Free Energy, G , of a state is: G = H – TS
 Free energy is a state function
 Therefore for a process at constant T, the free energy change will be:
G = H - TS
 The sign of G is important in predicting spontaneity
o If G 0, the forward reaction is spontaneous
o If G 0, the forward reaction is not spontaneous
 However the reverse reaction is
 Work must be supplied from the surroundings to make
the forward reaction go
o If G = 0, the reaction is at equilibrium
 The equilibrium position in a spontaneous reaction is given by the
minimum free energy available to the system
 Free energy decreases until it reaches a minimum value
Standard Free Energy Changes
 Standard states are pure solid, pure liquid, 1 atm(gas), 1 M
concentration(solutions) and 25C or 298 K
 Gf = 0 for elements
 The standard free energy change of a process is given by:
G =  nGf(products) - mGf(reactants)
Free Energy and Temperature Change
 The sign of G tells us if a reaction is spontaneous
 Focus on G =H- TS
o If H and -TS are less than zero, G must be less than zero
and the reaction is spontaneous
o If H and -TS are greater than zero, G must be greater than
zero and the reaction is not spontaneous
o I f  H and -TS have different signs, temperature will be an
important factor
 Thermodynamics gives us the direction not the rate of a reaction.
Free Energy and the Equilibrium Constant
 Recall that G and K(equilibrium constant) apply to standard
conditions
 Recall that G and Q apply to any conditions
 G = G + RTlnQ and at equilibrium Q=K and G =0
therefore G = -RTlnK
 We can conclude:
o If G 0, then K 1
o If G= 0, then K =1
o If G  0, the K 1