Chem Units 9-11

• Kinetics


  The study of reaction mechanisms and reaction rates
• Thermodynamics


  Does the reaction occur
• Reaction Mechanisms


  1. The step-wise process by which most reactions occur

  2. Few reactions occur in one step. Why?
• Example: production of carbon disulfide

  • Atoms collide for a moment to occur
  • Impossible for all to hit some time and energy (ping-pong ball example)
• Intermediates


  • Very unstable and short-lived
  • Produced in one step but lost in another step later in the reaction
  • Cannot isolate an intermediate
• Rate determining step


  • Slowest step in the reaction mechanism
  • To speed-up a reaction, you must speed-up the rate determining step
• Example: reaction between nitrogen dioxide and carbon monoxide to produce nitrogen monoxide and carbon dioxide

  1. Step #1: NO2 + NO

  2. Step #2: NO + CO
• Reaction rate


  The rate at which reactants disappear or products appear in a reaction (i.e. the speed of a reaction)
• Do all reactions occur at the same rate?
• Reaction rates vary from

  • Milliseconds to infinity
• Reasons reactions occur at different rates

  • The frequency of collisions
  • The effectiveness of collisions
• Effective collision


  A collision that results in a bond forming
• Conditions for an effective collision

  • The atoms collide at the proper angle
  • The atoms collide with the proper amount of energy
• Most collisions are not effective
• Collision theory


  The explanation for why reactions occur at different rates
• Factors affecting reaction rate

  • Nature of the reactants
  • Temperature
  • Concentration
  • Surface area
  • Catalyst
• Ionic compounds


  Reactions will be fast as the ions are free to react once dissolved in water
• Reaction of sodium carbonate and calcium chloride

  • Na₂CO₃ + CaCl₂ → 2NaCl + CaCO₃
• Molecular compounds


  Reactions will be slow as bonds must be broken in the reactants and reformed in the products
• As temperature increases


  The reaction rate increases
• Temperature effects


  • Hibernation and cooking
• As concentration increases


  The reaction rate increases
• Concentration effects


  • Zinc plus hydrochloric acid, lab experiments
• As surface area increases

  The reaction rate increases
• Surface area effects


  • Dissolving sugar, establishing lycopodium powder
• Catalyst


  Speeds up a reaction without being altered itself, lowers the activation energy of the reaction, provides a place for the reaction to occur
• Catalyst examples


  • Hydrogen peroxide and sodium potassium tartrate demo, lab, and catalytic converter
• Inhibitor


  Slows down a reaction
• When would you want to slow down a reaction?
• Thermodynamics


  The study of energy changes in a chemical reaction
• Enthalpy


  Also called the "heat of reaction", a measure of the difference in heat energy between the products and the reactants
• Standard heat of formation

  The energy released or absorbed when a mole of a substance is made from its elements
• Reference point: the elements in their natural state are defined as having 0 kJ/mol. The energy of a compound is based on this.
• Standard conditions for standard heat of formation: 1 mole of the substance is generated, temperature: 25°C, pressure: 1 atm, solution concentration: 1 M
• Calculating enthalpy changes


  • 2AgCl → 2Ag + Cl₂, ΔH=-121 kJ/mol

  4Al + 3O₂ → 2Al₂O₃, ΔH=-1670 kJ/mol
• Calculating ΔH


  • Na + 1/2 Cl₂ → NaCl, ΔH=-411 kJ/mol

  Na + 2H₂O → NaOH + 1/2 H₂, ΔH=-628 kJ
• Calculating the standard heat of formation of CaHO₃ given the enthalpy of the reaction CaO + CO₂ → CaCO₃ is -2194 kJ
• Spontaneous reaction


  Reactions that occur with little or no help from outside
• Nature tends to favor products that occur spontaneously
• Enthalpy


  Determines if a reaction is spontaneous