Chapter 16. Chemical Reactions Rates and Equilibria The rate of a chemical reaction shows how fast it goes. The equilibrium position of a chemical reaction shows how far it goes.

How Do Chemical Reactions Occur? Collision theory is a set of statements that explains how chemical reactions occur. It is an extension of the kinetic molecular theory. Kinetic molecular theory states that atoms, molecules, and ions, are in constant ran- dom motion. They collide with each other. Their average kinetic energy increases with temperature.

Collision Theory Reacting particles must collide with each other in order to react. Colliding particles must possess a minimum energy, called the activation energy, in order to react. Colliding particles must collide with the proper orientation in order to react.

Collision Theory NO 2 + CO  NO + CO 2

Collision Theory NO 2 + CO  NO + CO 2

Collision Theory

Why do Chemical Reactions Occur? Chemical reactions that are spontaneous occur to lower the energy of the system. They do this by maximizing favorable inter- actions between particles of opposite charge, while minimizing unfavorable interactions between particles of like charge.

Reaction Energy An endothermic reaction requires continuous energy input to continue. 6 CO H 2 O + light  C 6 H 12 O O 2  H  = 2816 kJ An exothermic reaction releases energy as it occurs. C 6 H 12 O O 2  6 CO H 2 O + heat  H  =  2816 kJ

Reaction Energy

The strength of the chemical bonds that form or are broken in a reaction determines whether it will be exothermic or endothermic. Breaking strong bonds absorbs more energy than breaking weak bonds. Forming strong bonds releases more energy than forming weak bonds.

Reaction Energy CH 3 OH + I 1-  CH 3 I + OH 1- Break CO350 kJ/mol Form CI 240 kJ/mol 110 kJ/mol For every mole of CH 3 OH that reacts, 110 kJ must be supplied to the reaction.

Reaction Energy CH 3 I + OH 1-  CH 3 OH + I 1- Break CI240 kJ/mol Form C O 350 kJ/mol 110 kJ/mol For every mole of CH 3 I that reacts, 110 kJ will be released.

Reaction Energy

Reaction Rates The rate of a reaction is the rate at which reactants are consumed or products are produced in a given time period. Depends on: Physical Nature of Reactants Reactant Concentration Reaction Temperature Presence of Catalysts

Reaction Rates Physical Nature of Reactants

Reaction Rates Reactant Concentration More particles mean more collisions!

Reaction Rates Reaction Temperature

Reaction Rates Catalysts

Chemical Equilibrium We've been writing chemical reactions as if they go to completion, or until at least one reactant is consumed. CH 3 OH + I 1-  CH 3 I + OH 1- But that only happens if the reaction releases a lot of energy. Usually things only go part of the way. CH 3 OH + I 1-  CH 3 I + OH 1-

Chemical Equilibrium In most chemical reactions, a mixture of products and reactants form, and at some point their concentrations stop changing. CH 3 OH + I 1-  CH 3 I + OH 1- There will still be a lot of CH 3 OH in the reaction mixture at that point.

Chemical Equilibrium Chemical equilibrium is the state reached when concentrations of reactants and products remain constant over time. Chemical equilibrium is reached when the rate of the forward reaction and the rate of the reverse reaction are equal. Consider the reaction H 2 (g) + I 2 (g)  2 H I (g)

Chemical Equilibrium

Equilibrium Stoichiometry mol NO and mol Br 2 are placed in a vessel. At equilibrium mol of NOBr is found. The reaction equation is: 2 NO(g) + Br 2 (g)  2 NOBr(g) Init mol mol  ?? ?? ?? Final ?? ?? mol

Equilibrium Stoichiometry mol NO and mol Br 2 are placed in a vessel. At equilibrium mol of NOBr is found. The reaction equation is: 2 NO(g) + Br 2 (g)  2 NOBr(g) Init mol mol  mol mol mol Final ?? ?? mol

Equilibrium Stoichiometry mol NO and mol Br 2 are placed in a vessel. At equilibrium mol of NOBr is found. The reaction equation is: 2 NO(g) + Br 2 (g)  2 NOBr(g) Init mol mol  mol mol mol Final mol mol mol

Equilibrium Constants The equilibrium constant is a numerical value that expresses the relationship between concentrations of reactants and concentrations of products for a system at chemical equilibrium.

Equilibrium Constants The equilibrium constant is calculated from the reaction equation and the equilibrium concentrations of reactants and products. 2 NO(g) + Br 2 (g)  2 NOBr(g) K eq = [NOBr] 2 = = 107 [NO] 2 x [Br 2 ] x

Equilibrium Constants The equilibrium expression is the symbolic form for the concentrations of reactants and products that is evaluated to obtain the equilibrium constant. a A + b B  c C + d D K eq = [C] c x [D] d [A] a x [B] b Concentrations of solids and pure liquids are 1 M.

Equilibrium Constants Example: 1.00 mole of acetic acid (CH 3 COOH) is dis- solved to make 1.00 L of aqueous solution. It ionizes according to the equation: CH 3 COOH + H 2 O  CH 3 COO 1- + H 3 O 1+ At equilibrium, [CH 3 COO 1- ] = [H 3 O 1+ ] = 4.2 x M What is K eq ?

Equilibrium Constants Example: Calcium carbonate is the business end of hard water. The maximum concentration of Ca 2+ and CO 3 2- ions in hard water is 7.07 x M. What is K eq for dissolving CaCO 3 in water, to form a saturated solution? CaCO 3 (s)  Ca 2+ (aq) + CO 3 2- (aq)

Equilibrium Constants Example: Lead(II) chloride is not very soluble in water. A saturated solution of PbCl 2 is M in Pb 2+ and M in Cl 1-. What is K eq for dissolving PbCl 2 ? PbCl 2 (s)  Pb 2+ (aq) + 2 Cl 1- (aq)

Equilibrium Constants So what do these things really tell me? If K eq is large (>1000), the concentrations of products are high relative to concentrations of reactants the position of equilibrium lies to the right the reaction is probably very exothermic

Equilibrium Constants If K eq is small (<0.001), the concentrations of reactants are high relative to concentrations of products the position of equilibrium lies to the left the reaction is probably very endothermic

Equilibrium Constants

If K eq is between and 1000, significant concentrations of both products and reactants are present at equilibrium.  H isn't large

Disturbing an Equilibrium One can shift the position of an equilibrium by applying a stress to the system. A stress to an equilibrium is a change in conditions, such as addition or removal of one of the components of the equilibrium, or changing its temperature. For gases, changing a concentration will involve changing a pressure.

Disturbing an Equilibrium LeChâtelier's Principle describes the behavior of an equilibrium which has been stressed. If a stress (change in conditions) is applied to a chemical system at equilibrium, the system will readjust (change the position of the equil- ibrium) in the direction that best reduces the applied stress.

Disturbing an Equilibrium 3 H 2 (g) + N 2 (g)  2 NH 3 (g) + heat What happens to the position of equilibrium if [N 2 ] is increased? [H 2 ] is decreased? The pressure is increased? [NH 3 ] is decreased? The system is heated? The system is cooled?

Disturbing an Equilibrium Being able to force a reaction that reaches equilibrium to completion is important. Separating the product(s) from reactant(s) takes time (time  $!) and disposing of waste reactants is expensive. C 7 H 6 O 3 + CH 3 COOH  C 9 H 8 O 4 + H 2 O Salicylic acid + acetic acid  aspirin + water K eq < 1, how do we get a high yield of aspirin?

Disturbing an Equilibrium What about catalysts? Catalysts effect rate, not equilibrium.

Disturbing an Equilibrium What about catalysts?