Chemistry 102 OUTLINE FOR KINETICS
I. RATES OF CHEMICAL REACTIONS
Definition Δ [M] /Δ time : The change in concentration of either product or reactant divided by the change in time.
Instantaneous : Measured at a single point using a tangent at that point. Average : Two point measurement graphing slope. Initial : Extrapolated to time zero! Used in most kinetic analyses
B. C. D.
II. ORDERS OF REACTIONS
A. First-order : A Z (where Z = one or more products)
1. Radioactive decay : Radioisotope breakdown
2. Half-life : t1/2 = 0.693/k (k = rate constant)
B. Second-order : A + B Z or 2A Z or 2B Z C. Third-order : A + B + C Z or 3A Z or 3B Z
III. FACTORS WHICH AFFECT RATES
A. B. C.
Molar Concentration of Reactant(s) Temperature Pressure
IV. RATE LAW
A. B. Definition : Rate = k [Reactant(s)]
How to Determine A Rate Law : See Handout with Examples. Rate Law may match stoichiometry of the reaction. Reaction has only one step! Rate Law may differ from the stoichiometry of the reaction. Reaction takes place in multiple steps.
FACTORS WHICH AFFECT REACTIONS
Activation Energy Orientation : Juxtaposition – exact positioning of reactants for productive collisions. Temperature Pressure Catalysts
C. D. E.
Definition : A series of steps by which a reaction takes place. The multiple steps must add up to the net chemical reaction.
Explanation : Some chemical reactions require more than one step. In many cases intermediates are formed which ultimately are used up and don’t appear in the net chemical reaction. Example : See Second Example on How To Determine a Rate Law Handout.
Definition: Substances that accelerate chemical reactions without being consumed.
Examples: Elements such as Ni & Pt. Biomolecules such as enzymes. How They Work: They generally lower the activation energy of a chemical reaction by increasing surface area and/or ordering molecules for productive collisions. Significance: They are everywhere.