Chemical kinetics: In what way do chemical reactions occur Important in medicine, cooking, industry Reaction mechanism (reaction pathway): Steps by which atoms rearrange themselves during a chemical reaction Can have several intermediate steps Reaction rate: How fast a reaction takes place Measured in moles of reactant disappearing or moles of product formed per unit time

Collision theory: In order for a reaction to occur, particles must collide with sufficient energy and the proper orientation Effective collision: a collision that has enough energy and the proper orientation, thus `resulting in a reaction taking place Activation energy (Ea): the minimum energy a collision must have to result in a reaction Rate of reaction depends on the number of effective collisions

Factors affecting rate of reaction: Nature of the reactants: The more tightly the atoms are bound, the slower the rate In solution, ionics react faster than covalents Gases react faster than liquids, which are faster than solids Simple molecules react faster than more complex ones

B) Concentration of reactants: The greater the concentration, the faster the rate of reaction The greater the concentration, the closer the particles are to each other, the greater the likelihood of collisions What about pressure? The higher the pressure, the closer the particles are to each other, the greater the concentration The higher the pressure, the faster the reaction

D) Temperature of a system: C) Surface area: The more surface area, the faster the rate Increases area of contact between reactants Ex: Steel bar versus steel wool Chewing of food D) Temperature of a system: The higher the temperature, the faster the rate Collisions are more likely to have sufficient energy to be effective

Catalyst speeds up reaction but is not consumed by the reaction E) Presence of a catalyst: Catalyst speeds up reaction but is not consumed by the reaction Lowers activation energy needed Provides alternate, lower energy pathway F) Presence of an inhibitor: Slows down reaction Ties up reactants Increases activation energy Ex: Preservatives, medicine

Energy in reactions (Thermodynamics) Heat: a form of energy; the ability to do work Potential energy (PE): stored energy Activation energy (Ea): Amount of energy for there to be an effective collision Amount of energy necessary to form activated complex Activated complex: A high energy, intermediate product Can continue to form final product

Heat of reaction (H, or enthalpy) H = heatproducts – heatreactants The amount of heat lost or gained during a chemical reaction (+) in endothermic reactions (-) in exothermic reactions H tells how endo or exothermic a reaction is

Potential energy (PE) diagrams: Reaction coordinate: How far a long a reaction has progressed Exothermic reaction: Gives off energy to surroundings So, heat of products is less than heat of reactants H is negative Heat is a product of the reaction Heat is written on right side of reaction arrow Ex: 2H2 + O2  2H2O + heat

Endothermic reaction: Absorbs energy from surroundings So, products have more energy than reactants H is positive Heat is a reactant Heat is written on the left side of reaction arrow Ex: 2H2O + heat  2H2 + O2

Effect of a catalyst on a chemical reaction: Lowers activation energy needed Provides alternate, lower energy pathway Speeds up the reaction in both directions Reaction just reaches equilibrium faster ΔH is unchanged

Chemical equilibrium: Dynamic equilibrium: The rate of one process is equal to the rate of the reverse process with no net change Types of equilibrium so far… Phase equilibrium: Melting = freezing or Evaporation = condensation B) Solution equilibrium: Dissolving = crystallizing

C) Chemical equilibrium: A condition in which… Rate of forward reaction = rate of reverse reaction *Must have a closed system Ex: N2 + 3H2 ⇌ 2NH3 + heat *The double-ended arrow indicates this is a reversible reaction At equilibrium, the rate at which N2 and H2 combine to form NH3 is equal to the rate at which NH3 breaks down into N2 and H2 There is no net change in the concentration or amount of N2, H2, or NH3

Lechâtelier’s Principle: If a stress is placed on a system at equilibrium, that equilibrium will shift to absorb the stress until a new equilibrium is reached Possible stress factors: 1) Increasing or decreasing concentration (Gases and aqueous) ONLY: Equilibrium will shift away from the side where concentration increases and toward the side where concentration decreases Ex: N2(g) + 3H2(g) ⇌ 2NH3(g) + heat Increase [H2], equilibrium shifts to the right Decrease [NH3], equilibrium shifts to the right

2) Increasing or decreasing pressure or volume (Gases ONLY): Increasing pressure will shift equilibrium to the side with the fewest moles of gas Decreasing pressure will shift equilibrium to the side with the more moles of gas Ex: N2(g) + 3H2(g) ⇌ 2NH3(g) + heat 1 mole 3 moles 2 moles 4 moles vs. 2 moles So, increasing pressure shifts equilibrium to the right

3) Increasing or decreasing temperature Treat “heat” as a reactant Increasing temperature shifts equilibrium away from side with “heat” on it Ex: N2(g) + 3H2(g) ⇌ 2NH3(g) + heat Increasing temperature shifts equilibrium to the left

OVERALL SUMMARY ON HOW TO PREDICT SHIFTS IN EQUILIBRIUM: 4) Presence of a catalyst: No change Rate of reaction increases equally in both directions Equilibrium is simply reached faster OVERALL SUMMARY ON HOW TO PREDICT SHIFTS IN EQUILIBRIUM: Equilibrium shifts away from side one adds to and toward the side one takes away from Changes in amounts of solids and liquids has no effect

Ex: H2(g) + Cl2(g) ⇌ 2HCI(g) + energy Ex: 2 ZnS(s) + 3 O2(g) ⇌ 2 ZnO(s) + 2 SO2(g) ∆H = +65Kcal Ex: C(s) ­+ H2O(g) ⇌ CO(g) + H2(g) Ex: 2 SO2(g) + O2(g) ⇌ 2 SO3(g) Ex: NH4CI(s) + heat ⇌ NH3(g) + HCI(g)

Spontaneous versus non-spontaneous reactions: Two basic natural tendencies: 1) Energy changes (H): Nature favors lower energy (-H) 2) Entropy changes (S): Entropy: randomness, disorder, chaos Nature favors higher entropy (+S) If a reaction has both +S and -H, it is always spontaneous If a reaction has both -S and +H, it is never spontaneous

Ways that entropy in a system can increase: Phase change solid liquid gas increasing entropy Ex: 2H2(g) + O2(g)  2H2O(l) indicates lower entropy B) Decomposition: Increase in the number of moles of particles Ex: 2H2O2  2H2O + 3O2 indicates higher entropy

C) Dissolving: Particles have more entropy in aqueous phase Ex: D) Number of moles of gas Ex: 2NH3(g)N2(g) + 3H2(g) 2 moles  4 moles, so an increase in entropy E) Temperature: Higher temperature means more randomness of particles, so an increase in entropy