Kinetics and Equilibrium Unit VI

I Kinetics A. Kinetics is the study of the rates of reactions and reaction mechanisms  Rate  Speed of a reaction  Mechanism  Steps involved in a reaction

B. Role of Energy in Kinetics 1. Activation Energy (E a )  Amount of energy needed to start up a reaction 2. Heat of Reaction (  H)  Measures the difference between the potential energy of the products and reactants in a reaction   H reaction = H products –H reactants  Reference Table I   H NO 2 = kJ   H NH 3 = kJ

C. Types of Reactions 1. Exothermic  Reactions that release energy  Energy is a product of the reaction  N 2 + 3H 2 → 2NH kJ   H is negative  “Feels hot” because it releases energy to you 2. Endothermic  Reactions that absorb energy  Energy is a reactant in the reaction  N 2 + O kJ → 2NO   H is positive  “Feels cold” because it absorbs energy from you

D. Collision Theory A reaction occurs when particles collide with sufficient energy and proper alignment More collisions, faster rate

E. Factors Affecting the Rate of Reactions 1. Concentration  Increasing concentration increases number of molecules  Increases collisions  Increases rate 2. Temperature  Increasing temperature increases speed of molecules  Increases collisions  Increases rate 3. Surface Area  Increasing surface area increases contact with reactants  Increases collisions  Increases rate

4.Pressure  Increase pressure decreases volume of a gas; less space  Increases collisions  Increases rate 5.Nature of Reactant  Ionic bonds break into pieces which increase concentration  Increases collisions  Increases rate  Covalent bonds do not breakdown ; react slower 6.Catalyst  Changes the mechanism of the reaction  Lowers activation energy making reaction occur easier  More collisions can occur  Rate increases

F. Potential Energy Diagrams 1. Exothermic Reaction 1 H reactant 2 activated complex 3 H product E a reverse E a forward  H

F. Potential Energy Diagrams 2. Endothermic Reaction 1 H reactant 2 activated complex 3 H product 4 E a forward 5  H 6 E a reverse

II Equilibrium Equilibrium is the balance between the rates of two opposing reactions Example of two opposing reactions: H 2 O(s) → H 2 O(l) melting H 2 O(s) ← H 2 O(l) freezing written in equilibrium notation H 2 O(s) ↔ H 2 O(l) phase change » forward direction—melt » reverse direction—freeze At equilibrium, Rates are equal Concentrations are constant

A. Types of Equilibrium Reactions 1.Phase Equilibrium -any phase changes 2.Solution Equilibrium -balance between dissolved and undissolved solute --saturated solution NaCl(s) ↔ NaCl(aq)  forward direction—dissolve  reverse direction—precipitate 3.Chemical Equilibrium -balance between forward and reverse reactions

B. Le Chatelier’s Principle systems that are at equilibrium are stable and want to remain at equilibrium when equilibrium reactions are “stressed”, they will “shift” to establish equilibrium stresses include: o concentration changes o temperature changes o pressure changes or volume changes (gases only) o add a catalyst

1. Concentration Changes When changing concentration, use the teeter- tauter technique Given the reaction: 2 NO 2 ↔ N 2 O 4 What will happen to [N 2 O 4 ] if [NO 2 ] is increased? ([ ] represents concentration) – Tilt left – Shift right – Makes more N 2 O 4

2. Temperature Changes When changing temperature, Use the teeter-tauter technique OR If temp increases, shift away from heat If temp decreases, shift toward heat Given the reaction: 2 NO kJ ↔ N 2 O 4 What will happen to [N 2 O 4 ] if the temperature is increased? Tilt left Shift right Makes more N 2 O 4 Heat (300 kJ) is on the left Equilibrium shifts away from the left side Shifts right N 2 O 4 increases

3. Pressure Changes Volume Changes When pressure changes occur, only gases will be effected (g) – Count gases in the system » An increase in pressure (decrease in volume) causes a shift toward the smallest side of the reaction Given the reaction: 2 NO 2 (g) ↔ N 2 O 4 (g) What will happen to [N 2 O 4 ] if the pressure is increased? – Count gases – 2 gases left side……….1 gas right side – Shifts right (fewer gases) – Makes more N 2 O 4

4. Addition of a Catalyst When a catalyst is added to a system at equilibrium, both forward and reverse reactions increase rate There is no effect on the equilibrium No shift will occur

Test yourself Given the reaction: N 2 (g) + 3 H 2 (g) ↔ 2 NH 3 (g) + 80 kJ What will happen to the concentration of NH 3 if H 2 increases? Temperature increases? Pressure increases? A catalyst is added?

III Enthalpy and Entropy There are TWO factors which determine if a reaction will occur spontaneously or not A. Enthalpy (ΔH) The natural tendency is to change to a lower energy state Exothermic direction is preferred B. Entropy (ΔS) Entropy measures randomness or disorder Greater disorder (messy), higher entropy  (solid) lowest entropy  (liquid)  (aqueous)  (gas) highest entropy High entropy is preferred

Spontaneous or Nonspontaneous? C (s) + O 2 (g) ↔CO 2 (g) + 120kJ ∆H = -120kJ Exothermic ∆ S solid/gas to gas only Increase in entropy Spontaneous

Spontaneous or Nonspontaneous? N 2 (g) + 4H 2 (g) + Cl 2 (g) + 93kJ ↔ 2NH 4 Cl (s) ∆H = +93kJ Endothermic ∆ S gas to solid Decrease in entropy Nonspontaneous

Spontaneous or Nonspontaneous? 2H 2 (g) + O 2 (g) ↔2H 2 O (l) + 80kJ ∆H = -80kJ Exothermic ∆ S gas to liquid Decrease in entropy ?????????? “Free Energy” needs to be used