Transcript Reaction Rates and Equilibrium

Reaction Rates and Equilibrium
M.Elizabeth
2011
Collision Theory
Used to Explain Reaction Rates
• Atoms, ions, and molecules can form a
chemical bond when they collide, provided
the particles have enough kinetic energy.
• Particles lacking the necessary kinetic energy
to react still collide, but simply bounce apart.
• Activation energy - the minimum energy
colliding particles must have in order to
react.
Chemical Reactions
• ordinarily occur as a result of collisions between
reacting particles. Consider the reaction:
• CO(g) + NO2(g) ----> CO2(g) + NO(g) rate = k(conc CO)(conc NO2(g))
▫
doubling the [CO], holding [NO2] constant, the number
of collisions in a given time doubles.
▫ doubling the [NO2] , holding CO constant, has the same
effect.
• the number of collisions per unit time is directly
proportional to the concentration of CO or NO2.
• The fact that the rate is directly proportional to these
concentrations indicates that reaction occurs as a direct
result of collisions between CO and NO2
molecules.
NOT EVERY COLLISION LEADS TO
REACTION!!!!!
• It is possible to calculate the rate at which
molecules collide with each other by using the
kinetic theory. Consider a mixture of CO and
NO2 at 700 K and a concentration of 0.10 mol/L
▫ every molecule would collide with about a billion
other molecules in one second!
▫ if every collision resulted in a reaction, then the
whole mixture would be reacted in a fraction of a
second.
▫ the actual reaction takes about 20 seconds.
Effective Collisions
• In order for collisions to be effective, there must
be considerable force in the collisions. The
slower moving molecules do not have enough
kinetic energy to react when they collide...they
bounce off one another and retain their identity.
• Only those molecules moving at high speed have
enough energy for collisions to result in a
reaction.
• Every reaction requires a certain minimum
energy for the reaction to occur--it is called
activation energy, Ea, and is expressed in kJ.
Activated Complex-an unstable,
high energy species • forward reaction
exothermic (∆H<0), Ea
is smaller than Ea1
• forward reaction
endothermic (∆ H>0),
Ea is larger than Ea1
• if ∆ H = +200 kJ, then
Ea = Ea1 + DH = Ea1 +
200 kJ
Factors that Affect Reaction Rates
Collision Theory
1. Temperature
2. Concentration
3. Particle size
4. Catalyst
Factors that Affect Reaction Rates
Collision Theory
Temperature
• Increasing temperature increases the
number of particles that have enough
kinetic energy to react when they collide.
Concentration changes (amt per vol)
• Cramming more particles into a fixed
volume increases the collision frequency.
Factors that Affect Reaction Rates
Collision Theory
Particle Size
• the smaller the particle size, the larger the
surface area for a given mass of particles.
Decreasing particle size will increase the
rate of reaction.
Catalyst
• A catalyst is a substance that increases the
rate of a reaction without being used up
itself in the reaction.
Effects of Catalyst on Activation
Energy
• Enzymes are
biological
catalyst usually
made of
proteins.
• Speed reactions
by lowering the
activation energy
of the reaction.
Chemical Equilibrium
Dynamic (in constant motion)
Reversible
• Chemical equilibrium occurs when the
forward and reverse reaction are taking
place at the same rate.
• There is no net change in the actual
amounts of the components of the system.
Equilibrium Example
Rate vs Equilibrium
At Equilibrium: RATES
ARE EQUAL
• the concentrations of
reactants and products
are constant.
D [ ]’s = 0
• The forward and
reverse reactions
continue after
equilibrium is
attained.
Kinetics Reaction Rate Orders
• the order of reaction with respect to a certain
reactant, is defined as the power to which its
concentration term in the rate equation is raised.
• For example, 2A + B → C
▫ r = k[A]2[B]1 the reaction order with respect to A
would be 2 and with respect to B would be 1, the total
reaction order would be 2 + 1 = 3.
• Reaction orders can be determined only by
experiment. The reaction order is not necessarily
related to the stoichiometry of the reaction, unless
the reaction is elementary. Complex reactions may or
may not have reaction orders equal to their
stoichiometric coefficients
The Reaction Quotient, Q
In general, all reacting chemical systems are
characterized by their REACTION QUOTIENT, Q.
When the
system is at
equilibrium,
Q=K
Reactions Review
• “Systems”: two reactions that differ only in
direction
• Any reversible reaction H2 + I2
noted by the double arrow;
↔ 2HI
↔
Two reactions: only difference is the Direction
H2 + I2 ↔ 2HI
Reactant
products
2HI ↔ H2 + I2
Left
Right
Reversible Reactions
H2 + I2 ↔ 2HI
• the products may react back to original reactants.
• “closed system”: ONLY if all reactant are present
• If one piece is completely gone it has ”gone to
competition” and no longer reversible
•
•
•
•
•
Examples: Reversible Reactions.
Unopened Soda
Breathing
Rechargeable batteries
Color changing shirt
The product is hexa-coordinated with water and the reactant is tetra-coordinated with water. Notice that the net charge on the left
and right sides of the first equation is zero.
Blue to pink
Co(H2O)4Cl2 + 2 H2O  Co(H2O)6Cl2
Properties of an Equilibrium
Pink to blue
Co(H2O)6Cl2 → Co(H2O)4Cl2 + 2 H2O
Blue to pink
Co(H2O)4Cl2 + 2 H2O → Co(H2O)6Cl2
Equilibrium systems are
DYNAMIC (in constant motion)
REVERSIBLE
can be approached from either
direction
Reaction at Equilibrium
A and B are _________. C and D are _______.
Factors Affecting Equilibrium
• Changes in temperature, pressure, and
concentration affect equilibrium.
• The outcome is governed by LE CHÂTELIER’S
PRINCIPLE
“...if a system at equilibrium is disturbed, the
system tends to shift its equilibrium position
to counter the effect of the disturbance.”
Writing and Manipulating Keq
Solids NEVER appear in equilibrium expressions.
S(s) + O2(g) ---> SO2(g)
Liquids NEVER appear in equilibrium
expressions.
NH3(aq) + H2O(liq) ---> NH4+(aq) + OH-(aq)
Henri Le Châtelier
• Henri Le Châtelier
• 1850-1936
• Studied mining
engineering.
• Interested in glass
and ceramics.
A +BC+D
Change in Concentration
• What happens when there is an increase
in reactant or product?
• What happens when there is a decrease
in reactant or product?
A +BC+D
Change in Concentration
Stress
Increase in A or B
Increase in C or D
Decrease in A or B
Decrease in C or D
Shift
forward
reverse
reverse
forward
Changes in Pressure
Only affects equilibrium system with
an unequal number of moles of
gaseous reactants and products
• Decrease in pressure shifts the reaction
in the direction that produces the larger
number of moles of gas.
• An increase in pressure shifts the
reaction in the direction that produces the
smaller number of moles of gas.
Product or Reactant Favored?
• Keq greater than 1 Products favored
• Keq less than 1 Reactants favored
Endothermic vs Exothermic
• Exothermic - Heat is a product. (-ΔH)
• Endothermic - Heat is a reactant. (+ΔH)
Changes in Temperature
• An increase in temperature favors
endothermic reactions.
• A decrease in temperature favors
exothermic reactions.
Changes in Temperature
• +ΔH = Endothermic
↑ temp favors the forward rxn.
• -ΔH = Exothermic
↑ temp favors the reverse rxn.
Temperature Effects on Equilibrium
N2O4 (colorless) + heat
2 NO2 (brown)
Equilibrium and Catalysts
• Add catalyst = no change
equilibrium concentration
• A catalyst only affects the RATE of
approach to equilibrium
Catalytic exhaust system.
Applying a Stress to a
System at Equilibrium
• N2 + 3H2  2NH3 + Heat
ΔH = -92 kJ/mol rxn
• Increase temperature
• Increase pressure
• Add a catalyst
• Adding H2
• Removing NH3
No Change
Work on Equilibrium Practice
Problems
The Equilibrium Constant
For any type of chemical equilibrium of the
type: a A + b B ↔ c C + d D
Keq is a CONSTANT (at a given Temp)
Equilibrium Constant
N2 (g) + O2 (g) ↔ 2NO (g)
Equilibrium Constant
CH4 (g) + Cl2 (g) ↔ CH3Cl (g) + HCl (g)
Magnitude of Keq
• Varies only with
temperature
• Is constant at a
given temperature
• Is independent of
the initial
concentrations
Magnitude of Keq
A+B↔C+D
• Keq > 1 mostly products
• Keq < 1 mostly reactants
• Keq ~ 1 equal amounts of
products and reactants
The Meaning of Keq
• Keq is used to tell if a reaction will favor
products or reactants.
• For N2(g) + 3 H2(g)  2 NH3(g)
• Concentration of products is much greater
than that of reactants at equilibrium.
• The reaction strongly favors products
The Meaning of Keq
• AgCl(s) ↔ Ag+(aq) + Cl (aq)
• Keq = 1.8 x 10-5
Concentration of
products is much less
than that of reactants at
equilibrium.
The reaction strongly
favors reactants
The Meaning of Keq
AgCl(s) ↔ Ag+(aq) + Cl (aq)
Keq = 1.8 x 10-5
Reactant Favored
The Reverse reaction
Ag+(aq) + Cl-(aq) ↔ AgCl(s)
Keq = 1.8 x 105
is product-favored.
Calculation of Keq – Learning Check
• For the reaction
2HI(g) ↔ H2(g) +I2(g), at 448•C,
• The equilibrium concentrations are
▫ HI =0.0040 M,
▫ H2 =0.0075 M, and
▫ I2 = 0.000043 M.
• Calculate the equilibrium constant at this
temperature.
▫ Write the equilibrium expression
(Prod/React)
▫ Fill in the values and solve
Calculation of Keq
• Equil
2HI(g) ↔ H2(g)
+ I2(g),
0.0040
0.0075
0.000043
Calculation of Keq Learning Check
N2 + 3H2  2NH3
• When equilibrium is established in a
5.00 L container the equilibrium
concentrations are: N2, 3.01 mol; H2,
2.10 mol; NH3, 0.565 mol.
• Calculate Keq
▫ Write the equation for Keq
▫ Calculate molar concentration (M)
▫ Fill in and solve.
Calculation of Keq
N2 + 3H2  2NH3
• [N2] = 3.01 mol/ 5.00L
• [H2] = 2.10 mol/ 5.00L
• [NH3] = .565 mol/ 5.00L
We must convert
concentrations to
Molarity = mol/L
• [N2] = .602 M
• [H2] = .420 M
• [NH3] = .113 M
Calculation of Keq
N2
• Equil
+ 3H2 ↔ 2NH3
0.602
0.420
0.113
Uses of Keq
We can us the
Keq to find
equilibrium
concentrations
at that
temperature.
Factors that Affect Equilibria
• Once a
reaction has
reached
equilibrium,
it remains at
equilibrium
until it is
disturbed by
some change
in conditions.