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 Activation Energy (Ea) Amount of energy needed to start up a reaction Heat of Reaction (H) Measures the difference between the energy of the products and reactants in a reaction Hreaction = Hproducts –Hreactants Reference Table I HNO2 = + 66.4 kJ HNH3 = - 91.8 kJ

Reactions that release energy C. Types of Reactions Exothermic Reactions that release energy Energy is a product of the reaction N2 + 3H2→ 2NH3 + 91.8kJ  H is negative ( H= - 91.8kJ) “Feels hot” because it releases energy to you Endothermic Reactions that absorb energy Energy is a reactant in the reaction N2 + O2 + 182.6kJ → 2NO  H is positive ( H= + 182.6kJ) “Feels cold” because it absorbs energy from you

D. Potential Energy Diagrams Exothermic Reaction Ea forward Ea reverse  H Hreactant HActivated Complex Hproduct

D. Potential Energy Diagrams Endothermic Reaction 1 Hreactant 2 Hactivated complex 3 Hproduct 4 Ea forward 5  H 6 Ea reverse

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

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

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

II Equilibrium Equilibrium is the balance between the rates of two opposing reactions Example of two opposing reactions: H2O(s) → H2O(l) melting H2O(s) ← H2O(l) freezing written in equilibrium notation H2O(s) ↔ H2O(l) phase change forward direction—melt reverse direction—freeze At equilibrium, Rates are equal Concentrations are constant

A. 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: concentration changes temperature changes pressure changes or volume changes (gases only) add a catalyst

1. Concentration Changes When changing concentration, use the teeter-tauter technique Given the reaction: 2 NO2 ↔ N2O4 What will happen to [N2O4 ] if [NO2 ] is increased? ([ ] represents concentration) Tilt left Shift right Makes more N2O4

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 NO2 + 300kJ ↔ N2O4 What will happen to [N2O4 ] if the temperature is increased? Tilt left Shift right Makes more N2O4 Heat (300 kJ) is on the left Equilibrium shifts away from the left side Shifts right N2O4 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 NO2 (g) ↔ N2O4 (g) What will happen to [N2O4 ] if the pressure is increased? Count gases 2 gases left side……….1 gas right side Shifts right (fewer gases) Makes more N2O4

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: N2(g) + 3 H2 (g) ↔ 2 NH3 (g) + 80 kJ What will happen to the concentration of NH3 if H2 decreases? Temperature increases? Pressure decreases? A catalyst is added?

III Enthalpy and Entropy There are TWO factors which determine if a reaction will occur spontaneously or not Enthalpy (ΔH) The natural tendency is to change to a lower energy state Exothermic direction is preferred 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) + O2(g) ↔CO2(g) + 120kJ ∆H = -120kJ Exothermic ∆ S solid/gas to gas only Increase in entropy Spontaneous

Spontaneous or Nonspontaneous? N2 (g) + 4H2(g) + Cl2(g) + 93kJ ↔ 2NH4Cl(s) ∆H = +93kJ Endothermic ∆ S gas to solid Decrease in entropy Nonspontaneous