Chem 102 Part One

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Chem 102 Part One
2013-02-28 09:23:40
Before first midterm

Before midterm
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  1. Kinetics
    Deals with the speed of reaction and its mechanisms
  2. Rate of Reaction
    Changes in concentration of reactants (or products) as a function of time

    • General Rate of reaction for:
    • aA + bB -> cC + dD
    • = -1/a(change in [A]/change in time)
  3. The Rate Law (or rate equation)
    • For a reaction relates the instantaneous reaction rate to the concentration of reactants
    • Only reactants concs appear in rate laws.

    • For aA + bB -> cC + dD differential rate law:
    • Rate = k[A]m[B]n
  4. Order
    Must be determined experimentally, generally not related to stoichiometry, typically integer (0,1,2) but can be negative or fractional
  5. Differential Rate Law (or rate equation)
    Relates instantaneous reaction rate to the concentration of reactants
  6. Integrated Rate Laws
    • Math expression giving reactant concentration as a function of time
    • -form depends on the order
  7. Reaction half-life
    • Time for reactant concentration to reach half its initial value
    • Characteristic of reaction under specific conditions
  8. Pseudo first order reaction
    • Higher order rxn (e.g., 2nd order) that appears like a first order rxn
    • Because the concentration of one (or more) reactants does not change significantly during the reaction
  9. Effect of Temp on Reaction rate
    • K=rate constant (for a given temp)
    • independent of reactant conc.
    • depends on temp or reaction

    • k=Ae-Ea/RT
    • Higher T, smaller exp term, Larger k

    For comparing two reactions to get Ea, don't forget to use Kelvins.  If you get a negative #, then you mixed up T1 and T2
  10. Collision Theory
    • Molecules must collide with the correct orientation and sufficient enery (>E a) to break and make bonds (to react)
    • A+B-> products
    • Concentration: Increasing concentration causes proportional increase in number of collisions
    • Not all collision cause a reaction
    • Only collisions with the right orientation will break/make bonds
    • Only collisions with enough energy to break/make bonds will cause reaction
  11. KMT (Kinetic Molecular Theory)
    says molecules have a distribution of kinetic energy
  12. Collision theory
    • To react moleucles must collide with the correct orientation and with sufficient energy to break and make bonds
    • Reaction involves:
    • reactants colliding
    • breaking old bond
    • making new bond
  13. Transition state
    • Unstable species formed by an effective collision between reactants
    • May relax to form either reactants or products
  14. Reaction Mechanisms
    • Most reactions occur via a series of small steps
    • Characteristics of these "elementary reactions":
    • involve just a single molecule (uni-molecular process) or two molecules (bi-molecular process)
    • Precise nature of each elementary process is postulated based on experimental results
    • Order of the rate law of an elementary process is the same as that step's stoichiometry
    • Can be reversible. ie. a step can reach equilibrium
  15. Reaction intermediates
    • Species that are formed in one step and then consumer in next step
    • Don't appear in the overall chemical equation or overall rate law
    • Formed, then consumed
  16. Rate determining step (rds)
    • Step in the rxn mech that controls rate of overall reaction
    • Slowest elementary reaction
    • Largest activation energy (Ea)
  17. Catalyst
    • Substance that speeds up rate of reaction without being consumer by the overall process
    • Does not appear in the overall chemical equation or overall rate law.
    • Lowers Ea
    • Speeds both fws and rev rxns
    • Provides new mech
    • Not consumed by overall reaction
  18. Pseudo first order reaction
    • Higher order rxn (eg 2nd order) that appears like a first order rxn
    • Because the conc of one (or more) reactants does not change significantly during the rxn
  19. Homogeneous Catalyst
    Substance in same solution as reactants and products that speeds up the rate of a reaction without being consumed by the overall process
  20. Heterogeneous Catalyst
    • Substance that is in a different phase than reactants and products that speeds up the rate of a reaction without being used up
    • Generally a solid
    • Mechanism involves 4 steps
    • Surface area of heterogeneous catalyst important
  21. Enzyme
    • Biological macromolecule (usually a protein) that acts as a catalyst
    • Increase rates by 108 or more
  22. Reaction Quotient (Q)
    Expresses the state of a system wrt products vs. reactants at a given moment

    Q = products/reactants

    • Q = K System at equilibrium
    • Q < K Excess reactants, system shifts to more products
    • Q < K Excess products, system shifts to more reantants
  23. For Pure Solids and Liquids
    • Pure solids and liquids (are at their standard state) do not have concentrations
    • Substitute 1 for their concentration in K and Q
  24. K
    • Is equilibrium
    • Little k is rate constant
    • Is unitless
    • Is constant at a given temp
    • Given initial concs, Q shows direction that the reaction will go to reach equilibrium
    • Q has the same form as K, except that it uses initial conc instead of equilibrium conc.
  25. ICE table
    • Initial
    • Change
    • Equilibrium
  26. Le Chatelier's Principle
    When a chemical system at equilibrium is disturbed, it retains equilibrium by undergoing a net rxn that reduces the effect of the disturbance
  27. Effect of change in concentration on equilibrium
    • cpd added -> equil shifts to consume added cpd
    • cpd removed -> equil shifts to replace cpd
    • solids/liquids added/removed -> no change
  28. Effect of change of pressure (volume) on equilibrium
    • adding/removing one gas component-> as discussed
    • adding inert gas -> no effect
    • changing volume -> possible change
    • increased volume: equilibrium shifts to higher number of total gas
    • decreases volume: shifts to less moles of gas
  29. Effect of Temperature on equilibrium
    • Need to know the thermodynamics of rxn
    • Treat heat as a reactant of product
    • Only conc and pressure appear in K
    • Exothermic: increase temp. K decreases
    • Endothermic: increase temp. K increases
  30. Arrhenius
    Acids and bases are compounds that dissociate to liberate H+ and OH-, respectively
  31. Bronsted-Lowry:
    • Acids want to donate a proton (H+) and bases want to accept a proton (H+)
    • If the acid donates a proton, a base must be present to accepts it.
  32. Strong acids
    • Hyrdohalic acids (HCl, HBr, HI)
    • Nitric acid (HNO3)
    • Perchloric acid (HClO4)
    • Sulfuric acid (H2HO4, Ka1)
  33. Strong Bases
    • LiOH, NaOh, KOH
    • Mg(OH)2, Ca(OH)2, Ba(OH)2
  34. Weak Acids
    • Acetic acids
    • Formic acids
    • Carboxylic acids
    • R-COOH
  35. Weak Bases
    • NH3
    • Methyl amine
    • Amine
    • R-NH
  36. Ka X Kb = Kw
  37. pKa + pKb = 14
  38. Buffer
    • Solution that lessens the impact of pH from addition of an acid or base
    • In pH range 3-11 mixture of acid and conj. base
    • A solution that resists changes in pH when acids or bases are added or when dilution occurs
  39. Buffer capacity
    • Measure of a solutions ability to resist pH change
    • Depends on: absolute concentration of components
    • relative concentration of components
  40. Buffer range
    • pH range over which a buffer is effective at resisting changes in pH when acid or base added
    • Buffer range = pKa +- 1
  41. Common ion effect
    • Shift in an equilibrium caused by addisiton of an ion involved in equilibiruim
    • Addition of common ion decreases solubitlity of a sparingly soluble salt
    • Removal of a common ion increases solubility of a sparingly soluble salt

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