1. Equilibrium aned kinetics

2. Essential Questions: What is chemical equilibrium
Essential Questions: What is chemical equilibrium? What are the types of chemical equilibrium?
Chemical equilibrium:
1. When the rates of the forward and reverse reactions are equal.
2. State of balance
3. No net change occurs in the actual amount of the compounds in the system.
Ex: 2SO2(g) +O2(g)  2SO3(g) 2:1:2
4. Dynamic state (both reactions continues), do not stop.
5. Represented by :
or 
6. Occurs only in a closed system

3. What is going on here? N2O4  2NO2
A FORWARD reaction goes from left to right (toward 2NO2).
A REVERSE reaction goes from right to left (toward N2O4)

4. It is a two-way reaction:
Observe the double arrow. The reaction is reversible.
N2O4(colorless)  2NO2(brown)
Molecules are constantly going back and forth between the two compounds.

5. Observe the reaction between H2 and I2:
1 mole H2 1 mole I2 2 moles of HI
The reaction is reversible. Observe the double arrow.

6. Reversible reaction-conversion of reactants to the products and conversion of the products to the reactants occurs simultaneously.
Ex: 2SO2(g) + O2(g)2SO3(g) forward reaction
2SO2(g) + O2(g)  2SO3(g) reverse reaction OR
2SO2(g) + O2(g) 2SO3(g) 

7. Over time, the concentrations reach the same levels regardless of the initial amounts.
At equilibrium, the amounts of products are constant (but not equal).

8. Shifting… C A D B A + B   C + D A B C D
Shift to the right (more products):
A B C D
Shift to the left (more reactants): A B
C D

9. An example of chemical equilibrium:
The Haber process for making ammonia:
N2 + 3H2  2NH kJ
Ammonia (NH3) is being made at the same rate as it is being turned back to N2 and H2.
Amounts of NH3, H2, and N2 remain constant.
Ammonia production critical to agriculture

10. Types of equilibrium: Phase Equilibrium Solution Equilibrium
1. Exists between the liquid and the solid phase.
2. Exists between the MP of the solids phase and the FP of the liquid phase Ex: H2O (s) H2O (L)

11. Solution Equilibrium:
1. When the rate of dissolving and the rate of recrystallization are equal.
2. Solution is saturated
Exs.: C12H22O11(s) C12 H22 O11(aq)
CO2(g) CO2(aq)

12. Equilibrium Position:
1. Relative concentration of the reactants and the products at equilibrium
2. Indicates whether the reactants or the products are favored.
Exs:
A  B Forward reaction favored formation of B
A  B Reverse reaction favors the formation of A.
3.Length of arrows indicates which reaction is favored.

13. Essential Questions: what is Le Chatelier’s Principle
Essential Questions: what is Le Chatelier’s Principle? What is the effect of different stresses on a system at equilibrium?
Le Chatelier’s Principle-if a stress is applied to a system in dynamic equilibrium, the system changes in a way that relieves the stress.
Stresses Include:
Concentration Product and Concentration Reactants
Change in Temperature
Change in Pressure
Catalyst

14. Concentration :
1. Change in concentration of the products or reactants disturbs equilibrium
2. System adjusts to minimize the effect of the change.
Exs. H2CO3(aq)CO2(aq) + H2O (L)
3. New equilibrium, will be established
When a reactant is removed, the reaction shifts in the direction of the formation of the reactants.
When a reactant is added, the reaction shifts in the direction of the formation of the products.

15. Temperature:
1. Increasing the temperature, causes the equilibrium to favor the endothermic reaction.
Ex: 2SO2(g) + O2 2SO3(g) + heat
2. Removal of heat favors the exothermic reaction
Ex: 2SO2(g) + O2 + heat 2 SO3(g)

16. Pressure:
1. Affects a system only in gaseous equilibrium that have an unequal number of moles of reactants and products.
2. Increase in pressure always favors the lower number of moles.
Ex: N2(g) + 3H2(g) 2NH3(g)
3. Decrease in pressure favors the higher number of moles.
Ex: N2(g) + 3H2(g)  2NH3(g)

17. Our example: N2 + 3H2  2NH3+ 92kJ
Effect on equilibrium
Endothermic rxn. Favored. Shift LEFT.
Exothermic reaction favored. Shift RIGHT.
Make more NH3 to compensate. Shift RIGHT.
Makes more N2 to compensate. Shift LEFT.
Makes more N2 and H2 to compensate. Shift LEFT.
favors NH3 (less space required). Shift RIGHT.
Favors N2 and H2 (more space allowed). Shift LEFT.
Stress:
Increase Temperature
Decrease temperature
Increase [N2]
Increase [H2]
Decrease [N2]
Increase [NH3]
Increase pressure
Decrease pressure

18. .
Common Ion – an ion that is found in both salts in a solution
Ex: PbCrO4(s) Pb+2 (aq) + CrO4-2(aq)
Addition of lead nitrate or sodium chromate, would cause a shift to the left to produce more PbCrO4

19. 2. Measured in amount of reactant changing per unit of time.
Essential Question: What are reaction rates, the collision theory and an activated complex?
Reaction Rates;
1. Measure of the spread of any change that occurs within an interval of time.
2. Measured in amount of reactant changing per unit of time.
Collision Theory:
1. Atoms/ions/ molecules can form products when they collide with each other if they have sufficient kinetic energy.
2. Particles without sufficient kinetic energy bounce apart unchanged.

20. Activation Energy-minimum energy that colliding particles must have in order to react.
Activated Complex:
1. Temporary intermediate product that may break apart and reform the reactants or rearrange that atoms and form new products the atoms and form new product.
2. Unstable arrangement of atoms
Forms at the peak of activation energy
4. In existence only briefly (10-13 seconds)

21. Potential Energy Diagrams –illustrate the potential energy that occurs during a chemical reaction

22. What is Energy of reaction? Energy gained or lost in a chemical change.
20 kJ are gained in this reaction.

23. Factors Affecting Reaction Rates: Temperature Concentration
Essential Questions: What are the four factors that affect the rate of a reaction? How does each factor affect the rate?
Factors Affecting Reaction Rates:
Temperature
Concentration
Particle Size
Use of Catalyst

24. Temperature:
1. Usually increasing temperature will increase rate of reactions.
2. Will increase the number of collisions and kinetic energy of the particles.
Concentration- An increase in the number of particles in a given volume, will increase the rate of the reaction. Ex: Crowded room

25. What does an endothermic reaction look like?
Energy is a reactant
A + B + energy  C + D
Energy is taken in from the surroundings.
Products contain more energy than the reactants.

26. What is Energy of reaction? Energy gained or lost in a chemical change.
20 kJ are gained in this reaction.

27. What does an Exothermic reaction look like?
Energy is a product
A + B  C + D + Energy
Energy given off to the surroundings.
Products contain less energy than the reactants.
Beaker gets hot or gives off light or heat.

28. Energy is lost in this chemical reaction. It exothermic.

29. Particle Size:
1. Total surface area of particles affects the rate of the reaction.
2. The smaller the particles, the larger the surface area for a given mass, thus an increase in the amount of reactants exposed for a reaction to occur.
Ex: Granulated sugar versus lump of sugar

30. Catalyst:
Substance that increase the rate of a reaction.
2. Permits reactions to occur at a lower energy level, thus decreasing the activation energy needed.
3. Not a reactant or a product
Ex: 2H2(g) + O2(g)  2H2O(L)
4. Unchanged in the reaction
Inhibitor-substance that interferes with the action of a catalyst

31. What does a catalyst do? It lowers the energy barrier.
← It is now easier to get over the hill. More particles have this amount of energy so more can react.

32. Kinetics – the study of how fast and how much

33. Essential Questions: What is entropy? How do you use Table I?
Enthalpy :
Heat of reaction
Measure in kJ at STP
3. Table I
In exothermic reactions , the potential energy of the products is lower than that of the reactants.
In endothermic reactions, the potential energy of the products is greater than that of the reactants.
Melting and evaporation are endothermic phases changes.
Condensation and freezing are exothermic phase changes. The greater the negative H , the more stable the products of the reaction.
The greater the positive H, the more unstable the products of the reaction.

34. A plot of Energy versus time for a chemical reaction:
DH = heat
Of reaction
Potential
Energy
PE of products
PE of Reactants
Progress of reaction 
In this case, energy is released (exothermic)

35. A plot of Energy versus time for a chemical reaction:
DH = heat
Of reaction
Potential
Energy
PE of products
PE of Reactants
Progress of reaction 
In this case, energy is absorbed (endothermic)

36. What is Heat of Reaction?
PE of products
Potential
Energy
DH = heat of reaction
DH= H(products) – H(reactants)
PE of Reactants
Progress of reaction 
Heat of reaction is the heat gained or lost in a reaction. It is the difference between Products’ Energy and Reactants’ Energy

37. What is activation energy?
ACTIVATED COMPLEX- highest energy reached.
ACTIVATION ENERGY is the energy required to get a reaction started.
It is the energy to get over the hill.

38. Where do I put the activation energy?
Potential
Energy
DH = heat Of reaction
DH= H(products) – H(reactants)
Progress of reaction 
Activation energy is the energy required to initiate (start) a reaction.

39. What is the effect of a catalyst?
Our catalyst is an enzyme↓
A catalyst lowers the energy hill. It lowers the ACTIVATION ENERGY.
The catalyst is not consumed in the reaction.

40. Reactions are reversible:
Activation energy for reverse (R L) reaction.
Activation
Energy
(forward)
Potential
Energy
DH= heat of reaction
Progress of reaction 
Activation energy is the energy required to initiate (start) a reaction.

41. Reference Table I

42. Entropy:
1. Measure of disorder of system
Ex: Clean room -low entropy
Messy room -high entropy
Solids-low entropy
Gases high entropy
2. Increase in entropy favors the spontaneous chemical reaction
3. Decrease in entropy favors the non-spontaneous reaction.

43. Law of Disorder- natural tendency for systems to move in the direction of maximum disorder or randomness
Spontaneous reaction:
1. Occurs naturally
2. Favors the formation of the products at the specified conditions
3. Releases energy.
Ex: fireworks

44. Nonspontaneous reactions-1
Nonspontaneous reactions-1.Doesn’t favor the formation of products under specified conditions. 2.
Ex: H2CO3(aq) CO2(g) + H2O(l) <1% >99%
99% of H2CO3 is converted to products in the forward reaction, thus it is spontaneous and releases energy
In the reverse reaction, CO2(g) + H2O(l), less that 1% combines to form H2CO3, thus nonspontaneous.
In most reversible reactions, 1 reaction is favors over the other

45. In every chemical reaction, heat is either release or absorbed and entropy or randomness either is increased or decreases. The size and direction of the enthalpy changes and entropy changes together determine whether a reaction is spontaneous; that is. Whether if it favors the products and releases energy.