1. Chemical Reaction:
Energy
Reaction Rates
Chemical Equilibrium

2. Chapter Outline Energy change in chemical reactions Reaction rate:
Concentration
Orientation
Activation energy
Catalyst
Chemical equilibrium: Le Châtelier’s principle:
Pressure/Volume
Temperature

3. Energy Energy do not have mass and volume (<-> Matter)
Energy is anything that has the capacity to do work
Chemistry is the study of matter, Chemical process involve energy change
it can cause physical and/or chemical changes in matter

4. Matter Possesses Energy
when a piece of matter possesses energy, it can give some or all of it to another object
all chemical and physical changes result in the matter changing energy

5. Kinds of Energy Kinetic and Potential
Kinetic Energy is energy of motion, or energy that is being transferred from one object to another (running water)
Potential Energy is energy that is stored (water above the dam)

6. Some Forms of Energy Electrical Heat or Thermal Energy
kinetic energy associated with the flow of electrical charge
Heat or Thermal Energy
kinetic energy associated with molecular motion
Light or Radiant Energy
kinetic energy associated with energy transitions in an atom
Nuclear
potential energy in the nucleus of atoms
Chemical
potential energy in the attachment of atoms or because of their position (example: Battery)

7. Units of Energy
calorie (cal) is the amount of energy needed to raise one gram of water by 1°C
kcal = energy needed to raise 1000 g of water 1°C
food Calories = kcals
Energy Conversion Factors
1 calorie (cal) =
4.184 joules (J)
1 Calorie (Cal)
1000 calories (cal)
1 kilowatt-hour (kWh)
3.60 x 106 joules (J)

8. Energy Use Unit joule (J) 4.18 3.6 x 105 9.0 x 108 calorie (cal) 1
Energy Required to Raise Temperature of 1 g of Water by 1°C °C
Energy Required to Light 100-W Bulb for 1 hr
Energy Used by Average U.S. Citizen in 1 day
joule (J)
4.18
3.6 x 105
9.0 x 108
calorie (cal) 1
8.6 x 104
2.15 x 108
Calorie (Cal)
0.001 86
215,000
kWh
1.1 x 10-6
0.10
250

9. Energy in Chemical Changes
Exothermic reaction: Reaction releases heat to the environment.
Product contains less energy than reactant.
Heat can be considered as PRODUCT.
Endothermic reaction: Reaction absorbs heat from the environment.
Product contains more energy.
Heat can be considered as REACTANT.

10. Examples of Exothermic Reaction
All Combustion reactions:
Hydrogen gas burns in the air:
O2 + 2H2  2H2O + heat
Acid-Base Neutralization Reaction:
NaOH + HCl  NaCl + 2H2O + heat
All explosive reactions

11. Examples of Endothermic Reaction
Mercury(II) oxide decomposes under heat
2HgO(s) + heat  O2(g) + 2 Hg(l)
Soda drink loses taste at room temperature:
H2CO3 + heat  CO2 + H2O

12. Bond Dissociation Energy
Breaking a stable chemical bond requires energy input: Endothermic reaction
Bond Dissociation Energy: energy to equally breaking a covalent bond.
Higher bond dissociation energy means stronger bond.

13. Calculate the Energy involved in Chemical Reaction
For a certain chemical reaction, the energy absorbed or released depends on the amount of product, or the limiting reactant.
The more product is generated, the more heat is involved.
The energy involved (kcal/mol or kJ/mol) can be considered as conversion factor.

14. Determine how much energy is released when 100
Determine how much energy is released when 100. g of propane is completely reacted with oxygen.
Problem solving plan:
Mass propane
mol propane
energy released
1.20x103 kcal

15. How Does Chemical Reaction Take Place?
In order for a reaction to occur, the reacting molecules must collide with each other.
Also two other factors:
Whether the collision has enough energy to “start to break the bonds holding reactant molecules together."
Whether the reacting molecules collide in the proper orientation for new bonds to form.

16. Effective Collisions
Effective Collisions: Collisions with sufficient energy and proper orientation (and therefore the reaction occurs).
The higher the frequency of effective collisions, the faster the reaction rate.
There is a minimum energy needed for a collision to be effective: Activation energy.

17. Effective Collisions: Kinetic Energy Factor
For a collision to lead to overcoming the energy barrier, the reacting molecules must have sufficient kinetic energy

18. Effective Collisions: Orientation Effect

19. Activation Energy
Reactants are often stable, requiring energy to re-to dissociate and form products.
Activation Energy: The energy barrier that prevents any collision between molecules from being an effective collision
The larger the activation energy of a reaction, the slower it will be.
Comparable to the Hurdles: The taller hurdle or slower runner, the harder for one to run overpass, fewer people will be able to clear.

20. Many Chemical Reactions Need Initial “Push” to Start
Rubbing a match head against a rough surface provides the activation energy needed for the match to ignite.

21. Reaction Energy Diagram
Transition state
(or Activated complex)

22. Exothermic Reaction Activation energy, large Activation energy, small
Reactants
Relative potential energy
DHreaction
Products
Progress of reaction

23. Endothermic Reaction Activation energy Products DHreaction Reactants
Relative potential energy
DHreaction
Reactants
Progress of reaction

24. Reaction Rates Some chemical reactions proceed rapidly.
Like the precipitation and acid-base reactions
Other reactions proceed slowly.
Like the discoloration of a sofa undersunlight.
The rate of a reaction is measured in the amount of reactant that changes into product in a given period of time.
Chemists study ways of controlling reaction rates.

25. Factors Effecting Reaction Rate: Reactant Concentration
The higher the concentration of reactant molecules, the faster the reaction will generally go. Why?
Increases the frequency of reactant molecule collisions
Examples:
High concentration (full strength) of CLORAX might overbleach or even damage fiber due to high reaction rate
When using Drano (NaOH-based solution), avoid running water (water dilutes NaOH)

26. Factors Effecting Reaction Rate: Temperature
Increasing the temperature, more molecules in the sample with enough energy to collide and overcome the activation energy.
Increasing the temperature also increases the frequency of collisions.
Example: Warm temperature quicken reaction such as CLORAX or DRANO in cleaning.

27. Catalysts
A catalyst is a substance that increases the rate of a reaction, but is not consumed in the reaction.
Catalysts lower the activation energy of a reaction by providing an easier pathway for the reaction.
Like lower the hurdles

28. Enzymes: Proteins as Biochemical Catalysts
Enzymes are protein molecules produced by living organisms that catalyze chemical reactions.
The enzyme molecules have an active site to which organic molecules bind.
 When the organic molecule is bound to the active site, certain bonds are weakened.
 Weakening bond makes chemical change easier.
i.e., the activation energy is lowered.

29. Amylase: An Enzyme Breaking down starch molecules

30. Overview of Factors Affecting Reaction Rates

31. Many Chemical Reactions are Reversible
If the products of a reaction are removed from the system as they are made, then a chemical reaction will proceed until the limiting reactants are used up.
However, if the products are allowed to accumulate, they will start reacting together to form the original reactants. Product  Reactant: Reverse Reaction.
Reactions that can proceed in both the forward and reverse directions are called Reversible Reactions.

32. Examples of Reversible Reactions
Photochromic sunglass: Transparent + Light Dark
Temperature dependent coloration:
N2O4 + heat 2NO2
Rechargeable battery:
Charged Depleted + electricity

33. Reaction Dynamics, Continued
The forward reaction _______ as the amounts of reactants decreases.
At the same time, the reverse reaction ________ as the concentration of the products increases.
Eventually, the forward reaction is using reactants and making products as fast as the reverse reaction is using products and making reactants. This is called Chemical Equilibrium.
Dynamic equilibrium is reached when the rates of two opposite processes are the same.

34. Equilibrium  Equal of forward rxn = of reverse rxn.
Concentrations of reactants and products are not always equal.
Some equilibria occur reactant ⇌ product
position of equilibrium favors the products.
Other equilibria occur reactant ⇌ product position of equilibrium favors the reactants.

35. Reaching Equilibrium: Population Changes
When country A feel overcrowded, some will emigrate to country B. The population of country B is so low that no one emigrates to counry A.
As time passes, emigration will occur in both directions at the same rate, leading to populations in country A and country B are constant, though not necessarily equal.

36. Equilibrium Constant
Even though the concentrations of reactants and products are not equal at equilibrium, there is a relationship between them.
For the reaction H2(g) + I2(g)  2HI(g) at equilibrium, the ratio of the concentrations in [ ] raised to the power of their coefficients is constant.

37. Equilibrium Constant, Continued
For the general equation aA + bB  cC + dD, the relationship is given below:
The lowercase letters represent the coefficients of the balanced chemical equation.
Always products over reactants.
The constant is called the equilibrium constant, Keq.

38. What Does the Value of Keq Imply?
When the value of Keq > > 1, we know that when the reaction reaches equilibrium, there will be many more product molecules present than reactant molecules.
The position of equilibrium favors products.
When the value of Keq < < 1, we know that when the reaction reaches equilibrium, there will be many more reactant molecules present than product molecules.
The position of equilibrium favors reactants.

39. A Large Equilibrium Constant: An Example

40. Disturbing and Re-Establishing Equilibrium
Once a reaction is at equilibrium, the concentrations of all the reactants and products remain constant.
However, if the conditions are changed, the concentrations of all the chemicals will change (not equilibrium) until equilibrium is re-established.
The new concentrations will be different, but the equilibrium constant will be the same.
Unless you change the temperature.

41. Le Châtelier’s Principle
Le Châtelier’s principle guides us in predicting the effect on the position of equilibrium when conditions change.
“When a chemical system at equilibrium is disturbed, the system shifts in a direction that will minimize the disturbance.”

42. Le Chatelier: Concentration Effect
Scenario A: [Reactant] , Other reactants  , products  (shift toward product)
That has the same Keq.
Scenario B: [Product] , Other products , reactants  (shift toward reactant)
You can keep removing product in an equilibrium to drive a reaction to completion!

43. Le Chatlier’s Principle: Population ChangesB A B
Since the population of A increases, rateforward < ratereverse , so the populations between A and B is no long at equilibrium, more people moving to B.
At equilibrium, rateforward = ratereverse ; so the populations stay constant. What if an influx of population enters A because of a war?
a new equilibrium between the populations is established. However, the new populations will have different numbers of people than the old ones. A B

44. Chemical Reactions cont’d
Concentration changes that result when H2 is added to an equilibrium mixture.

45. Practice—Predict the Effect on the Equilibrium When the Underlined Substance Is Added to the Following Systems:
2 CO2(g) Û 2 CO(g) + O2(g)
BaSO4(s) Û Ba2+(aq) + SO42-(aq)
CH4(g) + 2 O2(g) Û CO2(g) + 2 H2O(l)

46. Practice—Predict the Effect on the Equilibrium When the Underlined Substance Is Removed from the Following Systems:
2 CO2(g) Û 2 CO(g) + O2(g)
BaSO4(s) Û Ba2+(aq) + SO42-(aq)
CH4(g) + 2 O2(g) Û CO2(g) + 2 H2O(l)

47. Effect of Volume Change on Equilibrium
For solids, liquids, or solutions, changing the size of the container has no effect on the concentration.
Changing the volume of a container changes the concentration of a gas.
Same number of moles, but different number of liters, resulting in a different molarity.

48. Effect of Volume Change for Gas on Equilibrium
Decreasing the size of the container increases the concentration of all the gases in the container.
According to Le Châtelier’s principle, the equilibrium should shift to remove that pressure.
The way to reduce the pressure is to reduce the number of molecules in the container.
When the volume decreases, the equilibrium shifts to the side with fewer molecules.

49. The Effect of Volume Change on Equilibrium, Continued
Since there are more gas molecules on the reactants side of the reaction, when the pressure is increased the position of equilibrium shifts toward the products.
When the pressure is decreased by increasing the volume, the position of equilibrium shifts toward the side with the greater number of molecules—the reactant side.

50. Volume Changes on Equilibrium
Volume decrease: dark brown color fades away (2NO2→N2O4)
Volume increase: brown color turns darker (N2O4→2NO2)

51. Practice—Predict the Effect on the Equilibrium When the Volume Is Reduced.
2 CO2(g) Û 2 CO(g) + O2(g)
BaSO4(s) Û Ba2+(aq) + SO42-(aq)
CH4(g) + 2 O2(g) Û CO2(g) + 2 H2O(l)

52. The Effect of Temperature Changes on Equilibrium
Exothermic reactions release Heat:
Reactant  Product + Heat
Endothermic reactions absorb Heat:
Reactant + Heat  Product
Increase temperature = Add Heat
Decrease temperature = Remove Heat

53. Temperature on Equilibrium
Online video:
N2O4(colorless) Û 2NO2(brown): endothermic
Lower temperature favors N2O4

54. Temperature affects the Equilibrium
CoCl H2O Û Co(H2O) Cl- : exothermic
Lower temperature favors Co(H2O)62+

55. Practice—Predict the Effect on the Equilibrium When the Underlined Substance Is Added to the Following Systems, Continued:
2 CO2(g) Û 2 CO(g) + O2(g)
BaSO4(s) Û Ba2+(aq) + SO42-(aq)
CH4(g) + 2 O2(g) Û CO2(g) + 2 H2O(l)
Shift right, removing some of the added CO2 and
increasing the concentrations of CO and O2.
Shift left, removing some of the added Ba2+ and
reducing the concentration of SO42-.
Shift right, removing some of the added CO2 and decreasing the O2,
while increasing the concentration of CO2.

56. Practice—Predict the Effect on the Equilibrium When the Volume Is Reduced, Continued.
2 CO2(g) Û 2 CO(g) + O2(g)
BaSO4(s) Û Ba2+(aq) + SO42-(aq)
CH4(g) + 2 O2(g) Û CO2(g) + 2 H2O(l)
Shift left because there are fewer gas molecules
on the reactant side than on the product side.
No effect because none of the substances are gases.
Shift right because there are fewer gas molecules
on the product side than on the reactant side.