1. CHEMICAL KINETICS  Principally interested in… The rate of a chemical reaction The factors that influence the rate The mechanism by which a reaction takes place Rate of what? Change in concentration of a reactant or product per unit time. rate =rate = change in concentration change in time  Rates can be positive, negative or zero.  Positive – concentration increasing  Negative – concentration decreasing  Zero – concentration is constant  How do we measure [A]?  Spectroscopically….color change!  Production of a gas…pressure change!

2. THE COLLISON THEORY OF REACTION RATES  Collision theory is a model that accounts for the observed characteristics of reaction rates. It states that for a reaction to occur…  Particles must collide – (only two particles may collide at one time)  Particles must have the correct geometry.  Collision must involve enough energy to produce the reaction; that is, the collision must equal or exceed the activation energy, E a.  Bottom line…all collisions do not result in reactions! FACTORS THAT AFFECT REACTION RATES  1. Nature of the reactants – some reactant molecules react in a hurry, others slowly.  Physical state  gasoline(l) vs. gasoline(g)  K 2 SO 4 (s) + Ba(NO 3 ) 2 (s)  NRX; will react in aqueous state  Chemical identity – what is reacting?  Generally ions of opposite charge react very rapidly.  Generally the more bonds between reacting atoms in a molecule, the slower the reaction rate. Strong bonds react much more slowly.

3. 2. Concentration of reactants  More reactants, more collisions -­‐ the possibility of a successful collision increases! 3. Temperature  Heat ‘em up, speed ‘em up!  The faster they move the more likely they are to collide. Temperature does not affect activation energy. However, more particles now have sufficient energy to overcome the activation energy. Therefore, a general increase in reaction rate with increasing temperature.  A general rule of thumb is that a 10  C increase in temperature will double the reaction rate. 4. Catalysts  Increase rate but are not used up; not part of the chemical reaction Catalysts change the rate by providing an alterative reaction mechanism with a different activation energy.  Positive catalysts – increase reaction rate….lowers E a O demo with the cobalt chloride  Negative catalysts – decrease reaction rate…raises E a o Food preservatives!

4.  5. Surface area of reactants – exposed surfaces affect speed.  Except for substances in the gaseous state or solutions, reactions occur at the boundary, or interface, between two phases.  The greater the surface area exposed, the greater chance of collisions between particles, hence, the reaction should proceed at a much faster rate  Grain fires & coal dust are very hazardous and explosive  Solutions have the largest exposure!  Inert gases do not affect reaction rate…do not appear in rate law (more later) CHEMICAL REACTION RATES  The speed of a reaction is expressed in terms of its “rate” – some measureable quantity is changing with time. rate =rate =  Can be written in terms of reactant(s) disappearance or product(s) appearance change in concentration change in time

5. Example – 2NO (g) + O 2 (g) → 2NO 2 (g) Rate (slope) is not constant, it changes with time. The instantaneous rate can be found by finding the slope of a line tangent to the curve (see right). Rates can be positive or negative…  Product always POSITIVE  Reactant always NEGATIVE RELATIVE RATES Read the balanced equation….divide by stoichiometric coefficient  NO disappears and NO 2 appear twice as fast as O 2 disappears  This is truly an equality statement! EXAMPLE – What are the relative rates of change for the following reaction? 4PH 3 (g)  P 4 (g) + 6H 2 (g)

6. DIFFERENTIAL RATE LAW A few things first… Reactions are reversible and the reverse reaction is important. When the rate of the forward reaction equals the rate of the reverse reaction the reaction is at EQUILIBRIUM Though important, the reverse reaction is complicated…so study reaction soon after mixing to minimize reverse reaction (method of initial rates) Rate law (or rate expression) Rate law (or rate expression) – expression which shows how the rate depends on the concentration of reactant(s)  Depends only on the concentrations of the reactants (reverse reaction ignored)  Example: 2NO 2 (g) → O 2 (g) + 2NO(g) has the rate law format…  k, rate constant – proportionality constant that is temperature dependent, concentration independent & must be determined experimentally  n, reactant order – (generally) a positive integer 0, 1, 2 & must be determined experimentally. NOT balancing coefficients (more below) The rate constant, k, is temperature dependent!

7.  For additional reactants n, m & o must be determined experimentally!  A + B + C  products rate = k[A] n [B] m [C] o  Order of reaction – sum of reactant orders (n+m+o) Reactant order (n, m,…)  Zero order (n = 0)  A change in reactant concentration [A] has no effect on the reaction rate  Not very common: rate = k[A] 0 or rate = k  First order (n = 1)  A reaction rate is directly proportional to the reactant concentration [A]  Doubling [A] doubles reaction rate  Very common: rate = k[A] 1 or rate = k[A]  Nuclear decay usually fit into this category  Second order (n = 2)  A reaction rate is quadrupled when reactant concentration [A] doubles  Reaction rate increase by a factor of 9 when [A] is tripled, etc.  Common, particularly in gas-­‐phase reactions. rate = k[A] 2  Fractional orders are rare, but do exist! I’m confused…which reaction rate?  For the reaction 2NO 2 (g)  O 2 (g) + 2NO(g) you could write two reaction rates for NO 2 …these are equivalent!

8. Determining the rate law (rate equation)  Two techniques – depends on what the problem gives you!  Differential rate law  Given reactant(s) concentration(s) and initial reaction rate  Use “table logic” and method of initial rates  Integrated rate law  Given reactant concentration and time.  Find linear graph  Use AP equation sheet Differential rate law  BIG IDEA – find rate law…rate = k[A] n [B] m  Use “table logic”  Compare two experiments where one reactant concentration is constant. See how other reactant effects initial rate EXAMPLE 1 – Find rate law and rate constant, k for the following reaction. NH 4 + (aq) + NO 2 -­‐ (aq)  N 2 (g) + 2H 2 O(l) Experiment [NH 4 + ][NO 2 -­‐ ] Initial rate (M/s) 10.1000.0050 1.35  10 -­‐7 20.1000.010 2.70  10 -­‐7 30.2000.010 5.40  10 -­‐7

9. EXAMPLE 2 – Find the rate law and rate constant, k for the following reaction: BrO 3 -­‐ (aq) + 5Br -­‐ (aq) + 6H + (aq)  3Br 2 (l) + 3H 2 O(l) EXAMPLE 3 – The reaction below was studied and the following data were obtained: I -­‐ (aq) + OCl -­‐ (aq)  IO -­‐ (aq) + Cl -­‐ (aq) Determine the rate law, value for the rate constant and rate for an experiment where both [I -­‐ ] and [OCl -­‐ ] is initially 0.15 M. Experiment [BrO 3 -­‐ ] [Br -­‐ ][H+][H+]Initial rate (M/s) 10.10 8.0  10 -­‐4 20.200.10 1.6  10 -­‐3 30.20 0.10 3.2  10 -­‐3 40.10 0.20 3.2  10 -­‐3 Experiment[I -­‐ ][OCl -­‐ ]Initial rate (M/s) 10.120.18 7.91  10 -­‐2 20.0600.18 3.95  10 -­‐2 30.0300.090 9.88  10 -­‐3 40.240.090 7.91  10 -­‐2