1. Thermodynamics

2. Thermodynamics
All chemical changes are accompanied by changes in energy and the
degree of disorder for the particles involved.

3. Thermodynamics
Thermodynamics – the study of matter and energy interactions
enthalpy
entropy
predicts whether a reaction will occur spontaneously or not.

4. Enthalpy (H) heat content of a substance
most natural reactions are exothermic
natural systems tend to go from a state of higher energy to a state of lower energy.
products have less energy than the reactants and therefore are more stable
a ball rolls down a hill spontaneously, but not up. The ball looses potential energy as it rolls downhill. At the bottom of the hill it has zero potential energy. This same idea can be applied to chemical reactions.
Since some endothermic reactions do occur spontaneously there is another component.

5. Entropy (S)
a measure of the degree of randomness of particles in a system.
disorder of a substance.
ex: what happens if you carefully place three layers of different colored marbles in a container and shake it?
The marbles are no longer in an ordered state. No matter how much you shake the box, the probability of getting back to the three layers of single colors is next to zero.

6. Entropy
Just think of your locker. It does not spontaneously neaten itself (that takes work), it spontaneously becomes chaotic (messy).
a solution increases S compared to the pure solute and solvent.
Your room is probably an example.

7. Entropy: States of Matter
Crystalline solids have a repeating pattern – high degree of order
liquids are more disordered than those of solids
gas particles, moving almost totally at random, have the largest entropy of all
At absolute zero entropy is zero.
a solution increases S compared to the pure solute and solvent.
Your room is probably an example.

8. Entropy
Entropy increases when a substance is divided into more parts. Think of a jigsaw puzzle. The puzzle in the box (many individual pieces) has much more entropy than the assembled puzzle (one “piece”).
a solution increases S compared to the pure solute and solvent.
Your room is probably an example.

9. Entropy
In a chemical equation there is more entropy on the side that has a greater number of pieces.
6CO H2O  C6H12O O2
More entropy in the reactants (12 molecules) and less entropy in the products
a solution increases S compared to the pure solute and solvent.
Your room is probably an example.

10. Entropy: A positive value means an increase in the degree of disorder.
melting ice
A negative value means a decrease in the degree of disorder.
water freezing
Complete entropy worksheet
a solution increases S compared to the pure solute and solvent.
Your room is probably an example.

11. Δ Sorxn = Δ Soproducts - Δ Soreactants
Entropy equation
Δ Sorxn = Δ Soproducts - Δ Soreactants
a solution increases S compared to the pure solute and solvent.
Your room is probably an example.

12. Entropy sample problem #1
What is Δ So for calcium hydroxide given that when calcium reacts with water to produce calcium hydroxide and hydrogen Δ Sorxn = 25.7J/mol*k
u: Δ SoCa(OH)2
k: get ΔSo values from a table
ΔSoCa= 41.6 J/molK
ΔSoH2O= 69.9 J/molK
ΔSoH2= 131 J/molK
ΔSorxn= 25.7 J/K
a solution increases S compared to the pure solute and solvent.
Your room is probably an example.

13. Entropy sample problem #1
Ca(s) + 2H2O (l)  Ca(OH)2(s) + H2(g)
Δ SoCa= 41.6 J/molK; Δ SoH2O= 69.9 J/molK;
Δ SoH2= 131 J/molK; Δ Sorxn= 25.7 J/K
p: Δ Sorxn = Δ Soproducts - Δ Soreactants
s: 25.7J/K = [(1mol Ca(OH)2 *X) +
(1mol H2*(131J/molK))]- [(1mole Ca* 41.6J/molK) +(2mole H2O)*(69.9J/molK)]
X = 76.1 J/molK
a solution increases S compared to the pure solute and solvent.
Your room is probably an example.

14. Entropy sample problem #2
What is the change in standard entropy for the following equation?
H2(g) + Cl2(g)  2HCl(g)
So for H2(g) = J/molK
So for Cl2(g) = J/molK
So for HCl(g) = J/molK
a solution increases S compared to the pure solute and solvent.
Your room is probably an example.

15. Entropy sample problem #2
H2(g) + Cl2(g)  2HCl(g)
U: ΔSo
K: So for H2(g) = J/molK; So for Cl2(g) = J/molK; So for HCl(g) = J/molK
P: ΔSo= ΔSo(prod.) – ΔSo(react)
ΔSo= 2(So HCl) – (SoH2(g) + SoCl2(g) )
S:ΔSo=2(186.7J/molK) – (130.6J/molK J/molK)
ΔSo= 19.8 J/mol-K
a solution increases S compared to the pure solute and solvent.
Your room is probably an example.

16. Gibbs Free energy (G) chemical potential of a substance
comparison of the changes of enthalpy and entropy during a chemical reaction.
predicts spontaneity.

17. Gibbs Free energy All spontaneous reactions move toward equilibrium.
If the enthalpy and entropy for a chemical reaction have the same sign, there will be some temperature at which ΔH and T*ΔS will be numerically equal and ΔG will be exactly zero – equilibrium

18. Gibbs Free energy
Suppose that the enthalpy tends to make the reaction spontaneous, but the entropy change tends to prevent a spontaneous reaction. Is the reaction spontaneous or not?
There is an equation that will allow us to predict whether a reaction will be spontaneous or not.

19. Gibbs Free energy ΔGo = ΔHo - T ΔSo ΔGo = Gibbs Free Energy
ΔHo = enthalpy
T = temperature (check units)
ΔSo = entropy

20. Gibbs Free energy ΔGo = ΔHo - T ΔSo
When ΔGo for a reaction is negative, the reaction will be SPONTANEOUS.
When ΔGo for a reaction is positive, the reaction will be NONSPONTANEOUS.
When ΔGo equals zero it is in equilibrium.

21. Gibbs Free energy: sample
What is the change in free energy for the eq:
H2(g) + Cl2(g)  2HCl(g)
U: ΔGo
K: ΔHo= kJ/mol; ΔSo=19.8J/molK
P: ΔGo = ΔHo - TΔSo
S: ΔGo= kJ/mol – (298K)(0.0198kJ/molK)
ΔGo = kJ/mol
reaction is spontaneous

22. Gibbs free energy
Complete the worksheet

23. spontaneous nonspontaneous
Gibbs Free energy
ΔHo
ΔSo
ΔGo
spontaneous nonspontaneous - + -
spontaneous
nonspontaneous +
undeterminable ?
undeterminable ?

24. ΔGorxn = ΔGoproducts - ΔGoreactants
Equations
ΔHorxn = ΔHoproducts - ΔHoreactants
ΔSorxn = ΔSoproducts - ΔSoreactants
ΔGorxn = ΔGoproducts - ΔGoreactants
ΔGo = ΔHo - T ΔSo

25. Chemical Kinetics
The area of chemistry that is concerned with reaction mechanisms and rates reaction

26. Reaction process
Reaction mechanism – the step-by-step sequence of reaction by which the overall chemical change occurs
The pathway that a reaction takes as the reactants change into the products.
Elementary reactions = single step
Complex reactions = two or more steps

27. Reaction Process C(s) + 2H2(g)  CH4(g) ΔHof = ? given:
C(s) + O2(g)  CO2(g) ΔHoc=-393.5kJ/mol
2H2(g) + 1 O2(g)  2 H2O(l) ΔHoc=-571.6kJ/mol
CO2(g) H2O(l)  CH4(g) + 2O2(g) ΔHoc=+890.8
C(s) H2(g)  CH4(g) ΔHof= -74.3kJ/mol

28. Reaction process
Collision theory – in order for reactants to react (change into products) they must collide or come into contact with each other.

29. Effective collisions
Enough energy to overcome the energy barrier to break the bonds of the reactant particles and to form the bonds of the product particles. (activation energy)

30. Effective collisions
Proper orientation – reactant particles must collide in the correct positions to break and reform bonds.

31. Reaction Process
Activated Complex – The temporary combination of reactant particles formed as they collide with each other with enough energy to overcome the energy barrier.

32. Reaction Process
Intermediates – species that appear in some steps but not in the net equation.
C(s) + 2H2(g)  CH4(g) Δ Hof = ?
given:
C(s) + O2(g)  CO2(g) Δ Hoc=-393.5kJ/mol
2H2(g) + 1 O2(g)  2 H2O(l) Δ Hoc=-571.6kJ/mol
CO2(g) H2O(l)  CH4(g) + 2O2(g) Δ Hoc=+890.8
C(s) H2(g)  CH4(g) Δ Hof= -74.3kJ/mol

33. Reaction Process

34. Activated Complex

35. Reaction rates
The number reactant particles that react to form product particles per unit of time
The change in concentration of reactants per unit time as a reaction proceeds.

36. Rate-Influencing Factors
Temperature – the higher the temp., the faster the particles move and the more likely the reactant particles will have enough energy to overcome the energy barrier to form products.
Rule of thumb – for every 10oC rise in temp. the rxn rate doubles.

37. Rate-Influencing Factors
Concentration – the higher the concentration the higher the reaction rate.
Particle size/surface area – the smaller the particle size, the larger the surface area. The increase in surface area makes it more likely that reactant particles will collide to react to form product particles

38. Rate-Influencing Factors
Catalyst – reduce the activation energy needed for a reaction to occur
Catalysts – a substance that changes the rate of a chemical reaction without itself being permanently consumed.

39. Rate-Influencing Factors

40. Rate-Influencing Factors

41. Rate-Influencing Factors