University of Louisiana at Lafayette Chapter 6 Lecture Outline Prepared by Andrea D. Leonard University of Louisiana at Lafayette Modified by Peggy Sleevi Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

Energy Energy -- the ability to do work. Potential energy -- stored energy Kinetic energy -- energy of motion. Law of conservation of energy -- the total energy in a system does not change. Energy cannot be created or destroyed.

Types of Energy Energy can be transformed from one form to another Kinetic Potential Chemical Thermal Electrical Nucler Solar

Energy Chemical bonds store potential energy. A compound with lower potential energy is more stable than a compound with higher potential energy. Reactions that form products having lower potential energy than the reactants are favored.

Energy Units Calorie (cal) -- amount of energy needed to raise the temperature of 1 g of water by 1 oC. Joule (J) is another unit of energy. 1 cal = 4.184 J 1,000 J = 1 kJ 1,000 cal = 1 kcal 1 kcal = 4.184 kJ

Energy Changes in Reactions Chemical Reaction: bonds are broken in the reactants new bonds are formed in the products. Bond breaking requires an input of energy. Bond formation releases energy.

Bond Energy To cleave this bond, H = +58 kcal/mol. To form this bond,

Energy Changes in Reactions H = heat of reaction or enthalpy change H > 0 Energy is absorbed Reaction is endothermic H < 0 Energy is released Reaction is exothermic

Bond Dissociation Energy H for breaking a covalent bond by equally dividing e− ’s between the two atoms.

Bond Dissociation Energy Bond dissociation energies are positive values, because bond breaking is endothermic (H > 0). H H H + H H = +104 kcal/mol Bond formation always has negative values, because bond formation is exothermic (H < 0). H + H H H H = −104 kcal/mol

Bond Dissociation Energy The stronger the bond, the higher its bond dissociation energy. Bond dissociation energies generally decrease going down within a group.

Bond Dissociation Energies for Some Common Bonds ΔH (kcal/mol) H – H +104 Br – Br +38 F – F +58 Cl – Cl +46 I – I 636 H – OH +119 H – F +136 H – Cl +103 H – Br +88 H – I +71

Energy Changes in Reactions Calculations Involving H Values H indicates the relative strength of the bonds broken and formed in a reaction. When H is negative: Exothermic More energy is released forming bonds than needed to break the bonds. The bonds formed in the products are stronger than the bonds broken in the reactants. CH4(g) + 2 O2(g) CO2(g) + 2 H2O(l) H = −213 kcal/mol Heat is released.

Energy Changes in Reactions Calculations Involving H Values When H is positive: Endothermic More energy is needed to break bonds than is released in the formation of new bonds. The bonds broken in the reactants are stronger than the bonds formed in the products. 6 CO2(g) + 6 H2O(l) C6H12O6(aq) + 6 O2(g) ΔH = +678 kcal/mol Heat is absorbed.

Energy Changes in Reactions Calculations Involving H Values

Energy Diagrams For a reaction to occur, two molecules must collide with enough kinetic energy to break bonds. This bond breaks This bond forms

Energy Diagrams The orientation of the two molecules must be correct as well.

Energy Diagrams Ea, the energy of activation, is the difference in energy between the reactants and the transition state.

Energy Diagrams The Ea is the minimum amount of energy that the reactants must possess for a reaction to occur. Ea is called the energy barrier and the height of the barrier determines the reaction rate.

Energy Diagrams When the Ea is high, few molecules have enough energy to cross the energy barrier, and the reaction is slow. When the Ea is low, many molecules have enough energy to cross the energy barrier, and the reaction is fast.

Energy Diagrams The difference in energy between the reactants and the products is the H. If H is negative, the reaction is exothermic:

Energy Diagrams If H is positive, the reaction is endothermic:

Reaction Rates Increasing the concentration of the reactants: increases the number of collisions increases the reaction rate Increasing the temperature: increases the kinetic energy of the molecules increases the reaction rate

Reaction Rates – Catalysts A catalyst is a substance that speeds up the rate of a reaction. A catalyst is recovered unchanged in a reaction; does not appear in the product.

Reaction Rates – Catalysts In the following reaction, Pd acts as a catalyst: Catalysts accelerate a reaction by lowering Ea without affecting H.

Reaction Rates – Catalysts The uncatalyzed reaction (higher Ea) is slower. The catalyzed reaction (lower Ea) is faster. H is the same for both reactions.

Reaction Rates – Biological Catalysts Enzymes (usually protein molecules) are biological catalysts held together in a very specific three-dimensional shape. The active site binds a reactant, which then under- goes a very specific reaction with an enhanced rate. The enzyme lactase converts the carbohydrate lactose into the two sugars glucose and galactose. People who lack adequate amounts of lactase suffer from abdominal cramping and diarrhea because they cannot digest lactose when it is ingested.

Equilibrium A reversible reaction can occur in either direction, from reactants to products or from products to reactants. The forward reaction proceeds to the right. CO(g) + H2O(g) CO2(g) + H2(g) The reverse reaction proceeds to the left. The system is at equilibrium when the rate of the forward reaction equals the rate of the reverse reaction. The net concentrations of reactants and products do not change at equilibrium.

Equilibrium forward reaction CO(g) + H2O(g) CO2(g) + H2(g) reverse reaction The system is at equilibrium when the rate of the forward reaction equals the rate of the reverse reaction. The net concentrations of reactants and products do not change at equilibrium.

Equilibrium Constant The relationship between the concentration of the products and the concentration of the reactants [moles per Liter] is the equilibrium constant, K. For the reaction: a A + b B c C + d D equilibrium constant [products] [reactants] [C]c [D]d = K = = [A]a [B]b

Equilibrium Constant For the following balanced chemical equation: N2(g) + O2(g) 2 NO(g) [NO]2 equilibrium constant = K = [N2] [O2] The coefficient becomes the exponent.

Magnitude of the Equilibrium Constant When K is much greater than 1 (K > 1): [products] The numerator is larger. [reactants] Equilibrium favors the products and lies to the right. When K is much less than 1 (K < 1): [products] The denominator is larger. [reactants] Equilibrium favors the reactants and lies to the left.

Magnitude of the Equilibrium Constant When K is around 1 (0.01 < K < 100): [products] Both are similar in magnitude. [reactants] Both reactants and products are present. For the reaction: 2 H2(g) + O2(g) 2 H2O(g) K = 2.9 x 1082 The product is favored because K > 1. The equilibrium lies to the right.

Equilibrium Equilibrium favors the products when they are lower in energy than the reactants. K > 1 and ΔH < 0

Equilibrium

Calculating the Equilibrium Constant HOW TO Calculate the Equilibrium Constant for a Reaction Calculate K for the reaction between the general reactants A2 and B2. The equilibrium concentrations are as follows: Example [A2] = 0.25 M [B2] = 0.25 M [AB] = 0.50 M A2 + B2 2 AB Write the expression for the equilibrium constant from the balanced equation. Step [1] [AB]2 K = [A2][B2]

Calculating the Equilibrium Constant HOW TO Calculate the Equilibrium Constant for a Reaction Substitute the given concentrations in the equilibrium expression and calculate K. Step [2] [AB]2 [0.50]2 K = = [A2][B2] [0.25][0.25] 0.25 = = 4.0 0.0625 Since the concentration is always reported in mol/L, these units are omitted during the calculation.

Le Châtelier’s Principle If a chemical system at equilibrium is disturbed or stressed, the system will react in a direction that counteracts the disturbance or relieves the stress. Some of the possible disturbances: concentration changes temperature changes pressure changes

Le Châtelier’s Principle Concentration Changes 2 CO(g) + O2(g) 2 CO2(g) What happens if [CO(g)] is increased? The concentration of O2(g) will decrease. The concentration of CO2(g) will increase.

Le Châtelier’s Principle Concentration Changes 2 CO(g) + O2(g) 2 CO2(g) What happens if [CO2(g)] is increased? The concentration of CO(g) will increase. The concentration of O2(g) will increase.

Le Châtelier’s Principle Concentration Changes What happens if a product is removed? The concentration of ethanol will decrease. The concentration of the other product (C2H4) will increase.

Le Châtelier’s Principle Temperature Changes When the temperature is increased, the reaction that absorbs heat is favored. An endothermic reaction absorbs heat, so increasing the temperature favors the forward reaction.

Le Châtelier’s Principle Temperature Changes An exothermic reaction releases heat, so increasing the temperature favors the reverse reaction. Conversely, when the temperature is decreased, the reaction that adds heat is favored.

Le Châtelier’s Principle Pressure Changes When pressure increases, equilibrium shifts in the direction that decreases the number of moles in order to decrease pressure.

Le Châtelier’s Principle Pressure Changes When pressure decreases, equilibrium shifts in the direction that increases the number of moles in order to increase pressure.

Le Châtelier’s Principle Summary