Energy, Rates, and Equilibrium

1
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!

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 ClCl bond is measured to
  be 242 kJ/mol.

10
The enthalpy change required to break a
particular bond in one mole of gaseous molecules
is the bond energy.
Bond Energy
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 CH bond and one ClCl bond are broken
• one CCl and one HCl bond are formed.
• So,
• ?Hrxn D(CH) D(ClCl) ? D(CCl) D(HCl)
• (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 10C 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 427C, the
  following equilibrium is reached
• the equilibrium concentrations are I2 0.42
  mol/L, H2 0.025 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
LeChateliers Principle

• LeChateliers 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
LeChateliers 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
LeChateliers 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
LeChateliers 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
LeChateliers 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
LeChateliers 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
LeChateliers Principle

• summary of the effects of change of temperature
  on a system in equilibrium

43
Chapter 7
End Chapter 7