Reaction Rate How Fast Does the Reaction Go

Collision Theory l In order to react molecules and atoms must touch each other. l They must hit each other hard enough to react. l Anything that increase these things will make the reaction faster.

Energy Reaction coordinate Reactants Products

Energy Reaction coordinate Reactants Products Activation Energy - Minimum energy to make the reaction happen

Energy Reaction coordinate Reactants Products Activated Complex or Transition State

Energy Reaction coordinate Reactants Products Overall energy change

Things that Effect Rate l Temperature l Higher temperature faster particles. l More and harder collisions. l Faster Reactions. l Concentration l More concentrated closer together the molecules. l Collide more often. l Faster reaction.

Things that Effect Rate l Particle size l Molecules can only collide at the surface. l Smaller particles bigger surface area. l Smaller particles faster reaction. l Smallest possible is molecules or ions. l Dissolving speeds up reactions. l Getting two solids to react with each other is slow.

Things that Effect Rate l Catalysts- substances that speed up a reaction without being used up.(enzyme). l Speeds up reaction by giving the reaction a new path. l The new path has a lower activation energy. l More molecules have this energy. l The reaction goes faster. l Inhibitor- a substance that blocks a catalyst.

Energy Reaction coordinate Reactants Products

Pt surface HHHH HHHH l Hydrogen bonds to surface of metal. l Break H-H bonds Catalysts

Pt surface HHHH Catalysts C HH C HH

Pt surface HHHH Catalysts C HH C HH l The double bond breaks and bonds to the catalyst.

Pt surface HHHH Catalysts C HH C HH l The hydrogen atoms bond with the carbon

Pt surface H Catalysts C HH C HH HHH

Reaction Mechanism l Elementary reaction- a reaction that happens in a single step. l Reaction mechanism is a description of how the reaction really happens. l It is a series of elementary reactions. l The product of an elementary reaction is an intermediate. l An intermediate is a product that immediately gets used in the next reaction.

+ This reaction takes place in three steps

+ EaEa First step is fast Low activation energy

Second step is slow High activation energy + EaEa

+ EaEa Third step is fast Low activation energy

Second step is rate determining

Intermediates are present

Activated Complexes or Transition States

Mechanisms and rates l There is an activation energy for each elementary step. l Slowest step (rate determining) must have the highest activation energy.

Thermodynamics Will a reaction happen?

Energy l Substances tend react to achieve the lowest energy state. l Most chemical reactions are exothermic. l Doesn’t work for things like ice melting. l An ice cube must absorb heat to melt, but it melts anyway. Why?

Entropy l The degree of randomness or disorder. lSlS l The first law of thermodynamics. The energy of the universe is constant. l The second law of thermodynamics. The entropy of the universe increases in any change. l Drop a box of marbles. l Watch your room for a week.

Entropy Entropy of a solid Entropy of a liquid Entropy of a gas l A solid has an orderly arrangement. l A liquid has the molecules next to each other. l A gas has molecules moving all over the place.

Entropy increases when... l Reactions of solids produce gases or liquids, or liquids produce gases. l A substance is divided into parts -so reactions with more reactants than products have an increase in entropy. l the temperature is raised -because the random motion of the molecules is increased. l a substance is dissolved.

Entropy calculations l There are tables of standard entropy (pg 407 and the index). l Standard entropy is the entropy at 25ºC and 1 atm pressure. l Abbreviated Sº, measure in J/K. The change in entropy for a reaction is  Sº= Sº(Products)-Sº(Reactants). Calculate  Sº for this reaction CH 4 (g) + O 2 (g)  CO 2 (g) + H 2 O(g)

Spontaneity Will the reaction happen, and how can we make it?

Spontaneous reaction l Reactions that will happen. l Nonspontaneous reactions don’t. l Even if they do happen, we can’t say how fast. l Two factors influence. l Enthalpy (heat) and entropy(disorder).

Two Factors l Exothermic reactions tend to be spontaneous. negative  H. l Reactions where the entropy of the products is greater than reactants tend to be spontaneous. Positive  S. A change with positive  S and negative  H is always spontaneous. A change with negative  S and positive  H is never spontaneous.

Other Possibilities l Temperature affects entropy. l Higher temperature, higher entropy. l For an exothermic reaction with a decrease in entropy (like rusting). l Spontaneous at low temperature. l Nonspontaneous at high temperature. l Entropy driven.

Other Possibilities l An endothermic reaction with an increase in entropy like melting ice. l Spontaneous at high temperature. l Nonspontaneous at low temperature. l Enthalpy driven.

Gibbs Free Energy l The energy free to do work is the change in Gibbs free energy.  Gº =  Hº - T  Sº (T must be in Kelvin) l All spontaneous reactions release free energy. So  G <0 for a spontaneous reaction.

 G=  H-T  S HH SS Spontaneous? -+- At all Temperatures GG ++? At high temperatures, “entropy driven” --? At low temperatures, “enthalpy driven” +-+ Not at any temperature, Reverse is spontaneous

Problems l Using the information on page 407 and pg 190 determine if the following changes are spontaneous at 25ºC. 2H 2 S(g) + O 2 (g)  2H 2 O(l) + S(rhombic) l At what temperature does it become spontaneous?

2H 2 S(g) + O 2 (g)  2H 2 O(l) + 2S l From Pg. 190 we find  H f ° for each component – H 2 S = kJO 2 = 0 kJ – H 2 O = kJS = 0 kJ l Then Products - Reactants l  H =[2 (-285.8) - 0] - [2 (-20.1) + 1(0)] = kJ

2H 2 S(g) + O 2 (g)  2H 2 O(l) + 2 S l From Pg. 407 we find  S for each component – H 2 S = J/K O 2 = J/K – H 2 O = J/K S = 31.9 J/K l Then Products - Reactants l  S=[2 (69.94) - 2(31.9) ] - [2 (205.6) + 205] = J/K

2H 2 S(g) + O 2 (g)  2H 2 O(l) + 2 S l  G =  H - T  S l  G = kJ - 298K ( J/K) l  G = kJ J l  G = kJ kJ l  G = kJ l Spontaneous l Exergonic- it releases free energy. l At what temperature does it become spontaneous?

Spontaneous l It becomes spontaneous when  G = 0 l That’s where it changes from positive to negative. l Using 0 =  H - T  S and solving for T l 0 -  H = - T  S l -  H = -T  S l T =  H =  S = 1290 K kJ J/K = kJ J/K

There’s Another Way l There are tables of standard free energies of formation compounds.(pg 414)  Gº f is the free energy change in making a compound from its elements at 25º C and 1 atm. for an element  Gº f = 0 l Look them up.  Gº=  Gº f (products) -  Gº f (reactants) l Check the last problems.

2H 2 S(g) + O 2 (g)  2H 2 O(l) + 2S l From Pg. 414 we find  H f ° for each component – H 2 S = kJO 2 = 0 kJ – H 2 O = kJS = 0 kJ l Then Products - Reactants l  H =[2 (-237.2) - 2(0)] - [2 (-33.02) + 1(0)] = kJ

Reversible Reactions Reactions are spontaneous if  G is negative. If  G is positive the reaction happens in the opposite direction. 2H 2 (g) + O 2 (g)  2H 2 O(g) + energy 2H 2 O(g) + energy   H 2 (g) + O 2 (g) 2H 2 (g) + O 2 (g)  2H 2 O(g) + energy

Equilibrium l When I first put reactants together the for ward reaction starts. l Since there are no products there is no reverse reaction. l As the forward reaction proceeds the reactants are used up so the forward reaction slows. l The products build up, and the reverse reaction speeds up.

Equilibrium l Eventually you reach a point where the reverse reaction is going as fast as the forward reaction. l This is dynamic equilibrium. l The rate of the forward reaction is equal to the rate of the reverse reaction. l The concentration of products and reactants stays the same, but the reactions are still running.

Equilibrium l Equilibrium position- how much product and reactant there are at equilibrium. l Shown with the double arrow. l Reactants are favored l Products are favored l Catalysts speed up both the forward and reverse reactions so don’t affect equilibrium position.

Measuring equilibrium l At equilibrium the concentrations of products and reactants are constant. l We can write a constant that will tell us where the equilibrium position is. l K eq equilibrium constant l K eq = [Products] coefficients [Reactants] coefficients l Square brackets [ ] means concentration in molarity (moles/liter)

Writing Equilibrium Expressions l General equation aA + bB cC + dD l K eq = [C] c [D] d [A] a [B] b l Write the equilibrium expressions for the following reactions. l 3H 2 (g) + N 2 (g) 2NH 3 (g) l 2H 2 O(g) 2H 2 (g) + O 2 (g)

Calculating Equilibrium l K eq is the equilibrium constant, it is only effected by temperature. l Calculate the equilibrium constant for the following reaction. 3H 2 (g) + N 2 (g) 2NH 3 (g) if at 25ºC there 0.15 mol of N 2, 0.25 mol of NH 3, and 0.10 mol of H 2 in a 2.0 L container.

What it tells us l If K eq > 1 Products are favored l If K eq < 1 Reactants are favored

LeChâtelier’s Principle Regaining Equilibrium

LeChâtelier’s Principle l If something is changed in a system at equilibrium, the system will respond to relieve the stress. l Three types of stress are applied.

Changing Concentration l If you add reactants (or increase their concentration). l The forward reaction will speed up. l More product will form. l Equilibrium “Shifts to the right” Reactants  products

Changing Concentration l If you add products (or increase their concentration). l The reverse reaction will speed up. l More reactant will form. l Equilibrium “Shifts to the left” Reactants  products

Changing Concentration l If you remove products (or decrease their concentration). l The forward reaction will speed up. l More product will form. l Equilibrium “Shifts to the right” Reactants  products

Changing Concentration l If you remove reactants (or decrease their concentration). l The reverse reaction will speed up. l More reactant will form. l Equilibrium “Shifts to the left”. Reactants  products l Used to control how much yield you get from a chemical reaction.

Changing Temperature l Reactions either require or release heat. l Endothermic reactions go faster at higher temperature. l Exothermic go faster at lower temperatures. l All reversible reactions will be exothermic one way and endothermic the other.

Changing Temperature l As you raise the temperature the reaction proceeds in the endothermic direction. l As you lower the temperature the reaction proceeds in the exothermic direction. Reactants + heat  Products at high T Reactants + heat  Products at low T

Changes in Pressure l As the pressure increases the reaction will shift in the direction of the least gases. At high pressure 2H 2 (g) + O 2 (g)  2 H 2 O(g) At low pressure 2H 2 (g) + O 2 (g)  2 H 2 O(g)