R EACTION R ATES. R ATES OF C HEMICAL C HANGE Any measurable change in an activity expressed as a...

REACTION RATES

RATES OF CHEMICAL CHANGE Any measurable change in an activity

expressed as a function of time is a rate. Chemical Kinetics – the study of reaction

rates and the factors that affect them. CHEMICAL RATES MUST BE DETERMINED

EXPERIMENTALLY! Measureable changes for chemical reactions

are things like change in mass, volume of an evolved gas, change in the acidity of the solution, and change in color (concentration).

MEASURABLE CHANGES - EXAMPLE

RATES Concentrations must be measured often Reaction rates decrease with time (pendulum) RATES ARE POSITIVE, EVEN IF SUBSTANCE IS BEING

USED UP Rates for equilibrium reactions become constant

(forward rate = reverse rate) Rates for non- equilibrium reactions decrease to

zero (reaction goes to completion) The rate of change for each reactant or product

must be divided by its coefficient in the balanced chemical equation.

Rate = -1/a x Δ[reactant]/Δt = 1/b x Δ[product]/Δt a and b represent the coefficients of the balanced reaction The negative in front of 1/a is to keep the rate positive

MISCELLANEOUS RATE INFORMATION Ionic bonds are usually broken quickly

Covalent bonds are usually broken slowly

Homogeneous – all reactants are in the same phase

Heterogeneous – reactants are different phases

Rate = -1/a x Δ[reactant]/Δt = 1/b x Δ[product]/Δt2Br -(aq) + H2O2(aq) + 2H30+(aq) Br2(aq) + 4H2O(l)

2N2O5(s) 4NO2(g) + O2(g)

Time t (s) [H3O+] [Br2]

0 0.0500 085 0.0298 0.010195 0.0280 0.0110105 0.0263 0.0118

2Br -(aq) + H2O2(aq) + 2H30+(aq) Br2(aq) + 4H2O(l)

GRAPHICAL REPRESENTATIONS 1

H2(g) + I2(g) 2HI(g)

GRAPHICAL - RATE DECREASES WITH TIME

2N2O5(s) 4NO2(g) + O2(g)

CHEMICAL KINETICS – RATES AND MECHANISMS OF CHEMICAL REACTIONS Mechanisms – a step by step description of a

chemical reaction

Collision theory – particles must collide and must have proper angles and energy

FACTORS AFFECTING REACTION RATE Nature of reactants

Temperature of the system

Concentration of the reactants

Use of catalysts

Surface area

Pressure (partial pressure) is the same as concentration if reactants are all gases (Avogadro’s principle)

INTERMEDIATE PRODUCTS (ACTIVATED COMPLEX)

MECHANISMS All reaction mechanisms have a rate

determining step, a slow step.

TEMPERATURE Each 100C (10K) increase generally doubles

the rate, temperature increases the effectiveness of collisions

Higher temperature, more kinetic energy, more likely to reach the activated complex

COLLISIONS MUST HAVE PROPER ORIENTATION

CONCENTRATION Increase in concentration usually increases

the rate except for some homogeneous phase reactions or if that reactant is not involved in the mechanism

Higher concentration – greater frequency of collisions, more reactants with sufficient energy to reach the activated complex

CATALYSTS Catalyst – speeds up a reaction without being

altered itself, provides a new mechanism, lowers Eact

By the way, an inhibitor slows down a reaction (preservatives )

ENZYMES ARE SPECIFIC ORGANIC CATALYSTS

SURFACE AREA With greater surface area, reactants are

more likely to collide! Sugar cube vs. granular sugar

PARTIAL PRESSURES FOR GASES If the partial pressure of a gas increases, the

concentration increases (more gas, more pressure) and therefore the rate will also increases for reactions in which all reactants are gases.

HOW CAN REACTION RATE BE DESCRIBED?

RATE LAWS The rate law describes the way in which reactant

concentration affects reaction rates. General rate law rate ∝ [A]n = k [A]n The exponent (n) is the order of reaction with respect to A and must be determined experimentally. EXPONENTS ARE NOT simply THE COEFFICIENTS OF BALANCED REACTIONS AS WE DO WITH EQUILIBRIUM. (although they could be the same) 1ST order, n = 1, double the concentration, double the rate. 2nd order, n = 2, double the concentration, quadruple the rate. ETC. “0 “order, the reactant is not involved in the reaction rate. Simply “used up” as the reaction proceeds.

RATE LAW FOR MULTIPLE REACTANTSrate = k [A]n [B]mn is the order in [A]m is the order in [B]overall order = m + nUse experimental data to compare change in rate change in concentration.

DETERMINING RATE LAWS FROM DATA EXAMPLE 1

Experiment [Br2] [C3H60] Rate in M/s

1 .1 M .1 M 1.64 x 10-5

2 .2 M .1 M 1.65 x 10-5

3 .1 M .2 M 3.29 x 10-5

C3H60 + Br2 C3H5OBr + HBr

DETERMINING RATE LAWS FROM DATA EXAMPLE 2

Experiment [NO] [Br2 ] Rate in M/s

1 .003 M .003 M 4.04 x 10-4

2 .006 M .003 M 1.62 x 10-3

3 .003 M .006 M 8.08 x 10-4

2NO (g ) + Br2 (g ) 2NOBr(g )

ZERO ORDER GRAPHICALLY

1ST ORDER GRAPHICALLY

2ND ORDER GRAPHICALLY