1. Energy, Rates, and Equilibrium
Chapter 7
Energy, Rates, and Equilibrium

2. Energy
The capacity to do work

3. The Two Types of Energy
Potential: due to position or composition – One must compare what can be made from what you have in order to determine the potential.
Kinetic: due to motion of the object
KE = 1/2 mv2
(m = mass, v = velocity)

4. Food Calories Food is stored chemical potential energy.
Energy is produced by reacting the food we eat with oxygen.
Some of the energy can be used to do work!
We are all familiar with food and food energy. Read the nutritional information on any food product. Calories are everywhere. The study of thermodynamics has lead to our understanding of calorie content in foods.

5. Exothermic Reactions Some reactions give off heat.
Once the magnesium begins to burn, the heat given off makes more magnesium burn.
Link to Magnesium burning

6. Endothermic Reactions
Some reactions use heat during a reaction.
The hot water supplies the heat needed to sublime the solid carbon dioxide into gaseous carbon dioxide.
Link to Sublimation Video

7. Energy Potential - Stored energy
It is not the energy in the bonds of the compound you have in your hands that makes it have potential energy.
It is the energy of the new bonds that are formed in a chemical reaction that makes for the potential energy in the reactants.

8. Energy Diagrams
It is the difference which determines the potential

9. Covalent Bond Strength
Most simply, the strength of a bond is measured by determining how much energy is required to break the bond.
This is the bond enthalpy.
The bond enthalpy for a Cl—Cl bond is measured to be 242 kJ/mol.

10. Single bond < Double bond < Triple bond
The enthalpy change required to break a particular bond in one mole of gaseous molecules is the bond energy.
Bond Energy
H2 (g)
H (g) +
DH0 = kJ
Cl2 (g)
Cl (g) +
DH0 = kJ
HCl (g)
H (g) +
Cl (g)
DH0 = kJ
O2 (g)
O (g) +
DH0 = kJ O
N2 (g)
N (g) +
DH0 = kJ N
Bond Energies
Single bond < Double bond < Triple bond
9.10

11. Average Bond Enthalpies
This table lists the average bond enthalpies (kJ/mol) for many different types of bonds.
Average bond enthalpies are positive, because bond breaking is an endothermic process.

12. Energy
The energy required to break a bond is equal to the energy liberated when making a bond.
LAW of conservation of energy –
Neither created nor destroyed during physical or chemical changes

13. Enthalpies of Reaction
Compare the bond enthalpies of bonds broken to the bond enthalpies of the new bonds formed.
Hrxn = (bonds broken)(bonds formed)

14. Enthalpies of Reaction
CH4(g) + Cl2(g)  CH3Cl(g) + HCl(g)
In this example,
One C—H bond and one Cl—Cl bond are broken;
one C—Cl and one H—Cl bond are formed.
So,
Hrxn = [D(C—H) + D(Cl—Cl)  [D(C—Cl) + D(H—Cl)
= [(413 kJ) + (242 kJ)]  [(328 kJ) + (431 kJ)]
= (655 kJ)  (759 kJ)
= 104 kJ

15. Free Energy, Enthalpy, Entropy
DH = Total heat energy of a system at constant pressure.
Not all of this can be used to do work.
DG = That part which can do work
DS x T = That part which cannot do work.
DG = DH - T DS
If DG is favorable, the process will proceed and work can be extracted from the process.
If DG is unfavorable, work must be continually added to make the process proceed.

16. Reaction Rates
Even if energy can be extracted from a system, how fast the energy is produced can make or break the process literally.
Example: Atomic bomb vs. nuclear energy
Link

17. Reaction Rates
The rates of chemical reactions are affected by the following factors
molecular collisions
activation energy
nature of the reactants
concentration of the reactants
temperature
presence of a catalyst
On the following screens, we examine these factors one at a time

18. Molecular Collisions two species must collide to react
calculations show that the rate things collide is far greater than the rate at which they react
Conclusion: most collisions do not result in a reaction
a collision that does result in a reaction is called an effective collision
there are two main reasons why some collisions are effective and others are not;
energy
orientation

19. Molecular Collisions
Activation energy: the minimum energy required for a reaction to take place
energy is required for reactions to begin even if they give off energy during the process
this energy comes from collisions
if the collision energy is large, there is sufficient energy to break the necessary bonds, and reaction takes place
if the collision energy is too small, no reaction occurs

20. Molecular Collisions Orientation at the time of collision
the colliding particles must be properly oriented for bond breaking and bond making
for example, to be an effective collision between H2O and HCl, the oxygen of H2O must collide with the H of HCl so that the new O-H bond can form and the H-Cl bond can break

21. Energy Diagrams
Energy diagram for an exothermic reaction

22. Energy Diagrams
The reaction of H2 and N2 to form ammonia is exothermic
in this reaction, six covalent bonds are broken and six new ones are formed
breaking a bond requires energy, and forming a bond releases energy
in this reaction, the energy released in making the six new bonds is greater than the energy required to break the six original bonds; the reaction is exothermic

23. Energy Diagrams
Energy diagram for an endothermic reaction

24. Energy Diagrams Transition state: a maximum on an energy diagram
the transition state for the reaction between H2O and HCl probably looks like this, in which the new O-H bond is partially formed and the H-Cl bond is partially broken

25. Factors Affecting Rate
Nature of reactants
in general, reaction between ions in aqueous solution are very fast (activation energies are very low)
in general, reaction between covalent compounds, whether in water or another solvent, are slower (their activation energies are higher)
Concentration
in most cases, reaction rate increases when the concentration of either or both reactants increases
for many reactions, there is a direct relationship between concentration and reaction rate; when concentration doubles the rate doubles

26. Factors Affecting Rate
Temperature
in virtually all reactions, rate increases as temperature increases
an approximate rule for many reactions is that for a 10°C increase in temperature, the reaction rate doubles
when temperature increases, molecules move faster (have more kinetic energy), which means that they collide more frequently; more frequent collisions mean higher reaction rates
not only do molecules move faster at higher temperatures, but the fraction of molecules with energy equal to or greater than the activation energy also increases

27. Factors Affecting Rate
The distribution of kinetic energies (molecular velocities) at two temperatures

28. Factors Affecting Rate
Catalyst: a substance that increases the rate of a chemical reaction without itself being used up

29. Factors Affecting Rate
Many catalysts provide a surface on which reactants can meet
the reaction of ethylene with hydrogen is an exothermic reaction
if these two reagents are mixed, there is no visible reaction even over long periods of time
when they are mixed and shaken with a finely divided transition metal catalyst, such as Pd, Pt, or Ni, the reaction takes place readily at room temperature

30. Reversible Reactions
Reversible reaction: one that can be made to go in either direction
if we mix CO and H2O in the gas phase at high temperature, CO2 and H2 are formed
we can also make the reaction take place the other way by mixing CO2 and H2
the reaction is reversible, and we can discuss both a forward reaction and a reverse reaction

31. Reversible Reactions
Equilibrium: a dynamic state in which the rate of the forward reaction is equal to the rate of the reverse reaction
at equilibrium there is no change in concentration of either reactants or products
reaction, however, is still taking place; reactants are still being converted to products and products to reactants, but the rates of the two reactions are equal

32. Equilibrium Constants
Equilibrium constant, K: the product of the concentration of products of a chemical equilibrium divided by the concentration of reactants, each raised to the power equal to its coefficient in the balanced chemical equation
for the general reaction
the equilibrium constant is

33. Equilibrium Constants
Problem: write the equilibrium constant for this reversible reaction
solution: for this reaction, K is
note that no exponents are shown in this equilibrium constant; by convention the exponent “1” is not written

34. Equilibrium Constants
Problem: when H2 and I2 react at 427°C, the following equilibrium is reached
the equilibrium concentrations are [I2] = 0.42 mol/L, [H2] = mol/L, and [HI] = 0.76 mol/L. Using these values, calculate the value of K
Solution:
this K has no units because molarities cancel

35. Equilibrium and Rates
There is no relationship between a reaction rate and the value of K
reaction rate depends on the activation energy of the forward and reverse reactions; these rates determine how fast equilibrium is reached but not its position
it is possible to have a large K and a slow rate at which equilibrium is reached
it is also possible to have a small K and a fast rate at which equilibrium is reached
it is also possible to have any combination of K and rate in between these two extremes

36. LeChatelier’s Principle
LeChatelier’s Principle: when a stress is applied to a chemical system at equilibrium, the position of the equilibrium shifts in the direction to relieve the applied stress
We look at three types of stress that can be applied to a chemical equilibrium
addition of a reaction component
removal of a reaction component
change in temperature

37. LeChatelier’s Principle
Addition of a reaction component
suppose this reaction reaches equilibrium
suppose we now disturb the equilibrium by adding some acetic acid
the rate of the forward reaction increases and the concentrations of ethyl acetate and water increase
as this happens, the rate of the reverse reaction also increases
in time, the two rates will again become equal and a new equilibrium will be established

38. LeChatelier’s Principle
at the new equilibrium, the concentrations of reactants and products again become constant, but not the same as they were before the addition of acetic acid
the concentrations of ethyl acetate and water are now higher, and the concentration of ethanol is lower
the concentration of acetic acid is also higher, but not as high as it was immediately after we added the extra amount
the system has relieved the stress by increasing the components on the other side of the equilibrium
we say that the system has shifted to minimize the stress

39. LeChatelier’s Principle
Removal of a reaction component
removal of a component shifts the position of equilibrium to the side that produces more of the component that has been removed
suppose we remove ethyl acetate from this equilibrium
if ethyl acetate is removed, the position of equilibrium shifts to the right to produce more ethyl acetate and restore equilibrium
the effect of removing a component is the opposite of adding one

40. LeChatelier’s Principle
Problem: when acid rain attacks marble (calcium carbonate), the following equilibrium can be written
how does the fact that CO2 is a gas influence the equilibrium?
Solution: CO2 gas diffuses from the reaction site, and is removed from the equilibrium mixture; the equilibrium shifts to the right and the marble continues to erode

41. LeChatelier’s Principle
Change in temperature
the effect of a change in temperature on an equilibrium depends on whether the forward reaction is exothermic or endothermic
consider this exothermic reaction
we can look on heat as a product of the reaction
adding heat (increasing the temperature) pushes the equilibrium to the left
removing heat (decreasing the temperature) pushes the equilibrium to the right

42. LeChatelier’s Principle
summary of the effects of change of temperature on a system in equilibrium

43. Chapter 7
End
Chapter 7