Unit 8 Kinetics and Equilibrium

I. Kinetics  What does “kinetics” mean?  What do you think of when you hear kinetics?  A branch of chemistry that concerns itself with the speed of chemical reactions  Rates of reaction!

A. Collision Theory Collision Theory states: In order for reactants to change into products, the reactants must collide with the proper orientation and sufficient kinetic energy You must have an effective collision! alchemistry/flash/collis11.swf

B. Factors Affecting the Rate of a Reaction 1- Temperature: 2- Concentration: 3- Surface Area: 4- Bond Types: 5- Catalyst:  How do these factors increase or decrease the reaction rates???

Reaction Rates 1- Increases in Temperature cause an increase in the number of effective collisions, and this increases the rate of the reaction 2- Increases in concentration of reactants increase number of effective collisions, which increases the rate of reaction

Factors, (cont’) 3- Increase in surface area, essentially increases the concentration, gives more exposure to reactants, and increases effective collisions 4- Compatible bond types react faster than those without similar bond types/IMF’s

Last factor… 5- Addition of a catalyst will decrease the activation energy, therefore making the reaction proceed faster… Anytime the Ea is decreased, the reactants form the activated complex faster, and the reaction proceeds faster!

C. Potential Energy Diagrams  Drawings of how a reactant changes into a product  Show the energy changes associated with a chemical reaction Terms to Define: 1. Reactants 2. Products 3. Activated complex 4. Activation energy 5. Heat of reaction 6. Exothermic 7. Endothermic 8. Catalyst

Parts of the Potential Energy Curve

Potential Energy Diagram

Labeling a PE Diagram

Effects of a Catalyst  A catalyst will lower the activation energy of a reaction  Definition… [see review book!]  This will cause the reaction to proceed faster… –Why?  Online demo and references: agrams.htm

Catalyst Effects  Catalysts give reactants an alternative pathway to the intermediate, or activated, complex  This path requires less energy, therefore it becomes products faster!

Endothermic Reactions

Description of an Endothermic Reaction  PE of Products is GREATER than the PE or the reactants  (+) ΔHrxn  Needs heat energy to proceed!

Exothermic Reactions

Description of Exothermic Reactions  PE of the Products is LESS than PE of the reactants  (-)ΔHrxn  Releases energy!  chemistry/essentialchemistry/fl ash/activa2.swf chemistry/essentialchemistry/fl ash/activa2.swf

D. Heat of Reaction [ΔHreaction] Table I! Based on MOLE quantities!!! [look at coefficients within equation!] Negative ΔHrxn means…? Positive ΔHrxn means…?

II. Equilibrium  Equilibrium = a system that contains BOTH the reactants and products  The rate of the forward reaction is the same as the rate of the reverse reaction

A. Equilibrium Constants, Kc  Equilibrium constants are numerical values that compare the ratios of concentrations of products to the concentrations of the reactants  Contains gases and aqueous substances K C = [C] c [D] d [A] a [B] b

Kc Expression C For the reaction: aA + bB cC + dD where A, B, C and D are the chemical species and a, b, c and d are the stoichiometric coefficients, [C] c [D] d K 1 = [A] a [B] b

Descriptions of Kc  LARGE Kc value… suggests that the concentration of products present is much, much greater than the concentration of the reactants  Reaction went forward!!!!  SMALL Kc value… suggests that the concentration of reactants is very high and few products are made  Reaction did not go forward!

B. LeChatlier’s Principle 1. Definition: When a system is at equilibrium and it is disturbed, it will “shift” to correct the disturbance and reform a new equilibrium

LeChatlier’s Disturbances 1-Temperature Changes: An increase in Temperature will ALWAYS favor the endothermic reaction Ex] A + B C + energy Increases in Temperature similar to increasing ‘energy’; shifts to REVERSE direction!

2-Concentration Changes An increase in concentration for a reactant causes the reaction to shift to the forward direction to use the extra reactant An decrease in the concentration of a reactant causes the reaction to shift to the reverse direction to replace the missing reactant

3. Common Ion Effect  Adding a substance that generates the same ion already in the equilibrium system…  Increases the concentration of that ion!  Same effect as concentration!  Adding a substance that will bond with one of the ions in the system…  Decreases the concentration of that ion!  Same effect as concentration!

4. Pressure [for gases only]  An increase in the pressures of a gas will decrease the volume, therefore the system will shift to the side with the fewest number of moles of gas  A decrease in pressure causes the system to shift to the side with more moles of gas

Pressure effects cont’.  If a system has 1] the same number of moles of gas on each side OR 2] has no gases present  Changes in pressure HAVE NO EFFECT!!

5. Catalysts  If the system is AT equilibrium, the catalyst speeds BOTH the forward and reverse reaction rates  If the system has not yet REACHED equilibrium, the catalyst will increase the rate, causing it to reach equilibrium FASTER!

C. Spontaneity  A spontaneous reaction = one that occurs without the application of external work  Happens on its own!  Spontaneous in one direction, nonspontaneous in the opposite direction –Ex.] phase changes occur without help when the conditions are favorable!

Spontaneous Processes: [Enthalpy and Entropy]  Conditions favoring a spontaneous process: –These help explain why a process may proceed more easily… –Incorporates entropy changes from reactants to products  Formation of a gas  Formation of a precipitate  Formation of water