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CHM1051 > Kinetics > Flashcards

Kinetics Flashcards

1
Q

chemical kinetics

A

speed or rate of a reaction and its mechanism (do not confuse with the equilibria or extent)

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2
Q

rate

A

the change in concentration of reactants or products overtime (mol/L/s)

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3
Q

rate equation

A

rate = - 1/a (delta [A]/delta t) = -1/b (delta [B]/delta t) = 1/c (delta [C]/delta t) = 1/d delta [D]/delta t)

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4
Q

concentration and rate

A
  • substances must come into contact to react
  • more particles collisions, more reactions
  • rate usually increases with concentration
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5
Q

physical state and rate

A
  • for heterogeneous reactions (two different phases) collisions between reactants only occur at interfaces between phases
  • number of collision between the reactants per unit time depends on surface area of more condensed phase
  • increases surface area!!
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6
Q

temperature and rate

A
  • increase temperature, increases average kinetic energy of particles
  • as KE increases, particles move faster and collide more frequently with greater energy
  • increases rate
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7
Q

catalysts and rate

A

catalysts participate in a chemical reaction and increase the rate of the reaction withoutt undergoing a net chemical change itself

  • highly selective
  • often determine product by only speeding up one aspect of a reaction
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8
Q

average rate

A

the reaction rate between two points

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9
Q

instantaneous rate

A

using a tangent you can find the rate or gradient at any point

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10
Q

rate law

A

rate = - 1/a (delta [A]/delta t) = k [A]^n

rate: mol/L/sec
k is the rate constant and is characteristic of conditions
n is the rate order and is determined experimentally

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11
Q

first order

A

rate directly proportional to concentration of a reactant
k[A]^1

  • if k is doubled, [A] is doubled
  • if [A] decreases, rate decreases
  • plotted against time, it is non-linear
  • rate is fastest at beginning
  • k is s-1
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12
Q

second order

A

rate is proportional to the square of [A]
k[A]^2

if [A] doubles, rate quadruples
initially faster than first and then slows
k is L/mol/s

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13
Q

zero order

A

rate is independent of concentration
k[A]^0 = k

  • [reactant] graph is linear of -k
  • [product] is linear of +k
    molL/s
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14
Q

overall reaction order with more than one reactant

A

k[A]^n[B]^m

overall = n+m

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15
Q

determining rate order: method of initial rates

A
  • preform the reaction a number of times and vary the conditions (iodine clock)
  • look at if the rate doubles, quadruples etc. when concentration is changed
  • can determine k with algebra
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16
Q

determining rate order: integrated rate law

A

use if:

  • initial concentration of a reactant
  • concentration of a reactant overtime
  • several measures of [] between times
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17
Q

integrated rate law: zero order

A

[A]t = [A]0 -kt

linear [A] vs time

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18
Q

integrated rate law: first order

A

ln[A]t = ln[A]0 -kt
area under graph decreases linearly
linear if you plot ln[A] against t

19
Q

integrated rate law: second order

A

1/[A]t = 1/[A]0 + kt

linear if 1/[A] vs t

20
Q

half life for first order

A

half life: time it takes to halve a concentration to halve

t1/2 = ln2/k

  • independent of [A]
  • successive half lives do not change as [A] changes
21
Q

radioactive decay

A

the emission of a particle or photon that results from the spontaneous decomposition of the unstable nucleus of an atom

  • loss of a particle
  • loss of beta particle
  • emission of y radiation
  • rate is independent of chemical and physical form of the isotopes or temperatures
  • first order process
  • isotopes with shot half lives decay faster
22
Q

carbon dating

A
  • living things have constant C14:C12 ratio
  • C14 becomes N14 + beta- when they die
  • the half life is 5700 +/- 30 years
  • comparing ratio with that of living organisms helps determine age
23
Q

half life for second order

A

t1/2 = 1/k[A]0

24
Q

half life for zero order

A

t1/2 = [A]0/2k

25
Q

Arrhenius equation

A

k = Ae ^-(Ea/RT)

k is the rate constant 
A is the frequency factor 
Ea is activation energy 
R is 8.314J/k/mol
T is temp in k

ln(k2/k1) = -Ea/R(1/T2 - 1/T1)

26
Q

activation energy

A

the minimum energy required for a collision between molecules to result in a chemical reaction

  • energy to overcome electrostatic repulsion and a minimum amount of energy to break bonds so new ones can form
  • determines the rate of reaction
  • molecules that collide with less than Ea bounce off each other and are chemically uncharged (only direction and speed altered)
  • molecules with correct orientation (symmetry) react
27
Q

Boltzman distribution of KE

A
  • shaded areas (curve to baseline) are proportional to the total fractions of coliisons that involve the min Ea or more
  • at higher T, more molecules are present with sufficient KE for reactions - increases rate
  • area under curve is same for specified number of molecules
  • Ea is not temp dependent
28
Q

net reaction rate

A

sum of forwards and back reactions

rate = k1[A] - k-1[B]

29
Q

catalysis

A

alters the rate of a reaction without appearing in any of the products by providing a new energy pathway which may have different Ea

30
Q

homogenous catalysts

A

organometallic compounds and enzymes

  • catalyst is in same phase as reactants
  • collisions with reactants at ,ax because of uniform dispersion
  • no common mechanism
31
Q

heterogeneous catalysts

A
  • Haber process uses alpha-Fe catalyst
  • Zeolites for cracking petroleum
  • at least one of the reactants with the solid surface (adsorption) so that chemical bond in the reactant becomes weaker and breaks
32
Q

autocatalysis

A

one product catalyses the reaction eg. oxygen binds to haemoglobin

33
Q

homogenous

A

enzymes react in aqueous solutions within a cellular compartment

34
Q

heterogeneous

A

catalysts are embedded in the membrane

35
Q

enzymes are

A

specific

  • at low substrate concentrations, rate is first order with respect to [S]
  • at high concentrations rate is zero order with respect to [S]
36
Q

enzyme graph

A

steep increase and then flattens

  • V max is the fast, flat area
  • Km is half of Vmax

rate = V = (Vmax [S]/Km + [S])

37
Q

reaction mechanisms and elementary reactions

A

reactants to products (may be many steps)

  • can predict mechanism with elementary reactions
  • test prediction with experimentally determined kinetics data
38
Q

intermediates

A
  • produced in one step and used in another
  • do not appear in balanced equations
  • help determine rate
39
Q

rate determining steps

A
  • slowest elementary steps
  • orders of the rate equation are those from the rate determining step
  • have the highest Ea barrier
40
Q

molecularity

A

the number of molecules or atoms of reactant taking place

- unimolecular, bimolecular, trimolecular (rare) - harder to orientate

41
Q

unimolecular (A to products)

A

first order

42
Q

2A to products

A

bimolecular and second order

43
Q

A + B

A

bimolecular and second order

44
Q

2A + B

A + B + C

A

trimolecular and third

CHM1051 (9 decks)

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