R EACTION R ATES

R ATES OF C HEMICAL C HANGE 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).

M EASURABLE CHANGES - EXAMPLE

R ATES 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

M ISCELLANEOUS 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]/Δt 2Br - (aq) + H 2 O 2(aq) + 2H (aq)  Br 2(aq) + 4H 2 O (l) 2N 2 O 5(s)  4NO 2(g) + O 2(g)

Time t (s)[H 3 O + ][Br 2 ] Br - (aq) + H 2 O 2(aq) + 2H (aq)  Br 2(aq) + 4H 2 O (l)

G RAPHICAL REPRESENTATIONS 1 H 2(g) + I 2(g)  2HI (g)

G RAPHICAL - RATE DECREASES WITH TIME 2N 2 O 5(s)  4NO 2(g) + O 2(g)

C HEMICAL 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

F ACTORS 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)

I NTERMEDIATE PRODUCTS ( ACTIVATED COMPLEX )

MECHANISMS All reaction mechanisms have a rate determining step, a slow step.

TEMPERATURE Each 10 0 C (10K) increase generally doubles the rate, temperature increases the effectiveness of collisions Higher temperature, more kinetic energy, more likely to reach the activated complex

C OLLISIONS MUST HAVE PROPER ORIENTATION

C ONCENTRATION 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 E act By the way, an inhibitor slows down a reaction (preservatives )

E NZYMES ARE SPECIFIC ORGANIC CATALYSTS

S URFACE AREA With greater surface area, reactants are more likely to collide! Sugar cube vs. granular sugar

P ARTIAL 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.

H OW CAN REACTION RATE BE DESCRIBED ?

R ATE 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) 1 ST order, n = 1, double the concentration, double the rate. 2 nd 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.

R ATE LAW FOR MULTIPLE REACTANTS rate = k [A] n [B] m n is the order in [A] m is the order in [B] overall order = m + n Use experimental data to compare change in rate change in concentration.

D ETERMINING RATE LAWS FROM DATA EXAMPLE 1 Experiment [Br 2 ][C 3 H 6 0] Rate in M/s 1.1 M 1.64 x M.1 M1.65 x M.2 M3.29 x C 3 H Br 2  C 3 H 5 OBr + HBr

D ETERMINING RATE LAWS FROM DATA EXAMPLE 2 Experiment [NO][Br 2 ] Rate in M/s M 4.04 x M.003 M1.62 x M.006 M8.08 x NO (g ) + Br 2 (g )  2NOBr (g )

Z ERO ORDER GRAPHICALLY

1 ST ORDER GRAPHICALLY

2 ND ORDER GRAPHICALLY