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Chapters 15 & 16 (Equilibrium Systems (Dynamic Equilibrium is the…

- - - - ΔConcentration (adding/removing substances) Concentration increased - shift to opposite side Concentration decreased - shift to same side
      - ΔTemperature A system will shift to remove extra heat if added or make more heat if removed. Endothermic/positive enthalpy: cool-> reactants heat -> products Exothermic/negative enthalpy: cool -> products heat -> reactants
      - ΔPressure - affects only gaseous systems (inversely related to volume) Pressure increase - shift to less moles Pressure decrease - shift to more moles
      - Graph shapes
      - Catalysts
        
        Increase rate of reactions but do not affect concentration ratios or Kc
  - - - Change in State
        condensation/evaporation
    - - Solutes in solution
        crystallization/dissolution
    - - Blue Bottle reaction
        H₂ + I₂ ⇄ 2 HI
        forward/reverse
  - - - Heterogeneous substances can be left out of equilibrium equations
    - - Affected by: change in temperature and large changes in concentration
      - Communicates equilibrium concentration:
        Kc = 1 - system at equilibrium
        Kc > 1 - products favoured
        Kc < 1 - reactants favoured
- - - - Using data book: If acid is higher on table it favours the products. If the base is higher, it favours the reactants.
      - 1) List entities
        2) Identify acids and bases
        3) Identify strongest acid/base available
        4) Illustrate proton transfers
        5) Predict equilibrium position
    - - Acids donate a proton in the reaction, and bases accept protons.
      - Amphiprotic substances can accept or donate protons
    - - Pairs of substances on different sides of the equations that differ by a proton
  - - - Ka = Kw / Kb
      - Dissociation constant for weak acids into hydronium ions
      - Ka = [H3O+]^2/[acid equilibrium]
    - - Kb = Kw / Ka
      - Dissociation constant for weak bases into hydroxide ions
      - Kb = [OH-]^2/[base equilibrium]
    - - A Shortcut: Approximation Rule
        Can be used if [initial]/Ka or Kb > 1000 [H3O+] = √Ka x [acid initial]
        [OH-] = √Kb x [base initial]
      - Strong acids and bases ionize at 100%, so we replace them in equations with their respective H3O+ or OH- concentrations
  - - - Buffering region: pH changes very little
      - Equivalence point: equal number of moles of each reactant in solution
      - Buffering region/equivalence point #2: with polyprotic substances, it is possible a second buffering region and equivalence point can be produced for the second reaction
      - Buffer capacity: the limit of the buffer to maintain original pH
      - End point: After buffer capacity is reached, pH rises or lowers rapidly with addition of acid or base
    - - Buffers are combinations of a weak acid and its conjugate base having the same concentration in a solution.
      - The do NOT need to have a neutral pH of 7
      - Made up of mostly water, though there are significant amounts of the conjugate pair
      - pH of buffer changes very little when other acids or bases are added to the solution