Chapter 22

Reaction Rate & Chemical Equilibrium

Stability of Compounds

! In 2 TiO 2 Ti + O2 n  Overall energy change is (+)

w does not spontaneously decompose @ room temp.

n  Thermodynamically Stable

Stability of Compounds

! If overall energy change is (-), reaction will proceed spontaneously n  May be VERY slow n  C6H12O6 + 6O2 6CO2 + 6H2O

w @ room temp., no noticeable rxn w so slow it is Kinetically Stable

Stability of Compounds

∴ to predict whether a spont. rxn. will be useful, must know the rate @ which rxn. occurs and @ what pt. equilibrium is established.

Reversible Rxns. & Equilibrium

! Many rxns. result in an equilibrium mixture

! A rxn. goes to completion when all of one of the reactants is used up & rxn. stops n  Completion Rxn.

Reversible Rxns. & Equilibrium

! Completion Rxn. n  1 or more product is removed from rxn.

environment w gas is formed w PPT is formed w Water or undissociated, unionized subst. is

formed.

Reversible Rxns. & Equilibrium

! Not all rxns. go to completion H2(g) + I2(g) 2HI(g)

n  H2 & I2 make HI w bond betw. HI is weak & easily

decomposes to H2 & I2.

Reversible Rxns. & Equilibrium

! 1st rxn. goes from left to rt n  H2 + I2 2 HI

! 2nd rxn. goes from rt. to left n  H2 + I2 2 HI

! combined eqn. represents a reversible rxn. n  H2 + I2 2 HI n  Eventually reaches equilibrium

Reaction Rate

! If the product of a reversible rxn. decomposes faster than reactants form products, there will always be more reactant than product.

! Reaction Rate - the rate of appearance of a product or rate of disappearance of a reactant

Reaction Rate

n  usually units are (moles/ L) / s or M/s n  actually measures rate of change of

concentration ! If the 2 rxn. rates are known, we can

predict whether the product or reactant will be in higher concentration @ equilibrium.

Factors Affecting Reaction Rate

! Nature of reactants

! Concentration ! Temperature

! Catalysis ! Surface Area ! Pressure

n  gases only

Nature of Reactants

! Determines kind of rxn. that occurs n  Rxns. w/ bond rearrangement or e- transfer

take longer w neutral molec.

n  Ionic rxns. involve no e- transfer - faster n  Active metals & nonmetals react faster

than less active ones ∴ atomic structure affects rxn. rate

Nature of Reactants

! Formation of a new bond requires an “Effective Collision”

n  causes changes in e- clouds of colliding molecs.

n  Depends on: 1. Energy 2. Orientation

n  Colliding molecs. may form an Activated Complex w Unstable rxn. intermediate

Nature of Reactants

! Activation Energy - energy that must be attained in order for a collision betw. reactants to result in the formation of an activated complex n  energy to weaken or destroy original bonds n  If act. energy is high, few collisions have

enough energy to form activated complex w Very slow rxn w Kinetically stable

Concentration

! [ ] = mol / L - quantity of matter that exists in a unit vol. - molarity (M)

! For a rxn. to take place, particles must collide n  If # of particles per unit vol. (conc.) is incr.,

the chance of effective collisions is incr. n  If conc. of 1 reactant doubles, the rate may

double bec. twice as many collisions

Concentration

! Ex) A + B + C D n  If [A] is doubled, rate doubles n  If [A] & [B] are doubled, rate incr. 4X

! Ex) N2 + 3H2 2NH3 n  Rate1 = k1 [N2] - rate varies directly w/ [N2]

n  Rate2 = k2 [H2]3 - rate varies directly w/ [H2]

n  Rate3 = k3 [NH3]2

Concentration

! k is specific rate constant n  depends on size, speed, & kind of molecs

involved

n  ea. rxn. has only 1 value of k @ a given temp.

Concentration

! The rate expression for H2O2 + 2HI 2H2O + I2

is rate = k [H2O2] [HI] n  Even though 2 HI molecs. are in eqn., only

1 appears in the rate expression n  Only way to be sure of rate expression is to

use experimental data.

Concentration

! Rule of Thumb: n  Rxn. rate varies directly as the product of

the concen. of reactants w Not always true w To be sure, use experimental data

n  An incr. in press. on a gas will incr. its concen. & ∴ rxn. rate will incr.

Concentration

! Homogeneous rxn - reactants are all in the same phase

! Heterogeneous rxn. - rxn. which takes place @ the interface betw. 2 phases n  Ex) Zn dissolves (reacts) in H2SO4

w Rxn. takes place on the surface of Zn w ∴ if surface area is incr., rate of rxn. incr.

Concentration

! 2 H2 + O2 2 H2O ! Rate of formation = k[H2]2[O2] ! Find k if rate of formation = 0.6M/s; [H2] = 2.0 M; [O2] = 1.0M

Concentration

! In General for mA + nB C n  rate = k[A]m[B]n

w exponents are “order of the expression

n  Rate Laws are determined experimentally

Temperature

! Rxn. Rate is determined by frequency of collisions betw. molecs. n  If freq. of collisions incr., rate incr.

w for some rxns., their rate doubles for ea. 10 Co rise in temp.

Temperature

! An incr. in temp. will incr. K.E. of molecs. & ∴ collisions n  also incr. # of molecs. which have reached

activation energy ! An incr. in temp. will incr. the rate of rxn.

n  incr. # of activated complexes formed

Catalysis

! The process of increasing rxn. rates by the presence of a catalyst

! Catalyst - subst. which incr. a rxn. rate w/out being permanently changed n  decreases required activation energy

Catalysis

! Heterogeneous Catalyst - reactants & catalyst are not in the same state n  has a surface on which the substs. can

react. w adsorbs one of the reactants w Adsorbtion - the adherence of 1 subst. to the

surface of another n  ex) catalytic converters

Catalysis

! Homogeneous Catalyst - exists in same phase as reactants n  enters into the rxn. - forms rxn.

intermediate or activated complex w requires less activation energy

n  returns unchanged in final step of rxn.

Catalysis

! Inhibitors - “tie up” a reactant or catalyst in a complex so it will not react. n  does not slow down rxn. - stops it

Reaction Mechanism

! Most rxns. occur in a series of steps. n  usually involves collision of only 2 particles

w rarely involve 3 or more particles

Reaction Mechanism

! If a rxn. consists of several steps: A B; B C; C final product

One of the steps will be slower than all the others n  Rate Determining Step n  Faster steps will not affect the rate

Reaction Mechanism

! Reaction Mechanism - The series of steps that must occur for a rxn. to go to completion n  @ a given temp., the rate of a rxn. varies

directly w/ the product of the concentrations of the reactants in the slowest step.

Reaction Mechanism

2H2 + O2 2H2O ! Rate of formation = k [H2]2 [O2]

n  3rd Order ! A + B C R = k [A] [B]

n  2nd Order ! A + 2B C R = k [A] [B]2

n  3rd Order

Reaction Mechanism

N2 + 3 H2 2 NH3 R = k [N2] [H2]3 n  4th Order

! Sum of the exponents is the Order of the Expression

Reaction Mechanism

! If rxn. is a single step rxn., coef., in eqn. will become exponent in rate expression n  The only way to know the rate expression

for sure is by examining experimental data.

Equilibrium Constant

H2 + I2 2 HI (Forward rxn.) ! As rxn starts, lots of H2 & I2, no HI

n  as rxn. proceeds, there’s less & less H2 &I2 w fewer molecs. mean fewer collisions

n  There’s more & more HI w rxn. of 2HI H2 + I2 is incr. (reverse rxn.)

Equilibrium Constant

! When the rate of forward rxn. = rate of reverse rxn., we have equilibrium n  rate of forward rxn. = kf [H2] [I2] n  rate of reverse rxn. = kr [HI]2

! @ equilibrium: n  kf [H2] [I2] = kr [HI]2

Equilibrium Constant

! kf = constant kr ! Equilibrium Constant - Keq = kf

kr ! Solve for kf / kr ! Keq = [HI]2

[I2] [H2]

Equilibrium Constant

! General eqn n  for mA + nB sP + rQ n  Keq = [P]s [Q]r = [Prod.] [A]m [B]n [React]

Equilibrium Constant

! If Keq is small (<1), very little product is formed. n  Reactant is favored.

! If Keq is lg. (>1), rxn. is nearly complete n  much product is formed n  product is favored.

Equilibrium Constant

! What is the equilibrium constant for the following rxn. if the final concentrations are CH3COOH = 0.302M, CH3CH2OH = 0.428M, H2O = 0.654M, and CH3CH2OOCCH3 = 0.655M?

CH3COOH + CH3CH2OH H2O + CH3CH2OOCCH3

Equilibrium Constant

! What is the equilibrium concentration of SO3 in the following rxn. if the concentrations of SO2 and O2 are each 0.0500M and Keq = 85.0?

2SO2 + O2 2SO3

Le Chatelier’s Principle

! Conditions affecting equilibrium: 1. Temp. 2. Press. 3. Concentration (of prods. & reacts.)

! If a condition is changed (stress) on a syst. in equilib., then the equilib. will shift to restore the original conditions (relieve the stress).

Le Chatelier’s Principle

N2(g) + 3H2(g) 2NH3(g) + energy 1. Conc. of reactants is incr. (either H2 or

N2) n  # of collisions betw. reactants incr

n  Incr. rxn. rate toward right (shift right) n  amt. of product formed is incr.

Le Chatelier’s Principle

N2(g) + 3H2(g) 2NH3(g) + energy 2. Press. is incr.

n  Has same effect as incr. conc. of all gases in eqn.

n  Applies only to gases n  Equilib. usually shifts to right

w ck equilib. expression

Le Chatelier’s Principle

! Keq = [NH3]2 [N2] [H2]3

n  If press. doubles, reverse rxn. must speed up by a factor of 4

n  since [H2] is cubed doubling press. (which doubles conc.) speeds up forward rxn. by a factor of 16

Le Chatelier’s Principle

! In H2(g) + Cl2(g) 2HCl(g) n  Doubling press. will not shift equilib.

w Why? w Rate in ea. direction is affected the same way.

! An incr. is press. will always drive a rxn. in the direction of the smaller # of moles of gas. n  Press. affects only gases

Le Chatelier’s Principle

N2(g) + 3H2(g) 2NH3(g) + energy 3. If temp. is incr., equilib. may shift either

left or right. n  If heat is a product, equilib. will shift left n  If heat is a reactant, equilib. will shift right

Optimum Conditions

! Conditions which produce hightest yield.

In Haber process: 1. High conc. of H2 & N2 should be

maintained. 2. NH3 should be removed as it’s formed. 3. Temp. should be high enough to maintain

a reasonable rate, but low enough not to favor reverse rxn.

Optimum Conditions

4. Catalyst should be used to lower activation energy

5. High press. should be maintained.