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Chem 2 Heider > Ch 13 - Reactions > Flashcards

Ch 13 - Reactions Flashcards

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1
Q

ectotherms

A

organisms with their body temperature dependent on their surroundings

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2
Q

What is the effect of cold on reptiles?

A

the rate of reactions in their muscles occur more slowly making it impossible to move quickly

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3
Q

How does the rate of a reaction affect a rocket?

A

extremely important.

Too slow and no lift.

Too fast and can explode.

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4
Q

What does the rate of a chemical reaction measure?

A

How fast the reaction occurs

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5
Q

A fast reaction rate =

A

a larger fraction of the molecules react to form products in a given period of time.

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6
Q

H2(g) + I2(g) -> 2HI(g)

A
  • Rate = -(delta[H2]/(delt(t))) = -([H2]t2-[H2t1])/t2-t1
    - Rate = -(delta[I2])/(delta(t))
    - the hydrogen concentration at t1 and t2 divided by the change in time
    - the reaction rate is defined as the negative of the change in concentration divided by the change in time
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7
Q

Why is the reaction rate defined as negative?

A
  • it is negative because the initial concentration of the reactants decreases as the reaction proceeds
    - the change in the concentration of a reactant is negative
    - the negative sign in the equation makes the overall rate positive(this is by convention that reaction rates are reported as positive quantities)
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8
Q

Rate with respect to the product = +(1/2)(delta[HI]/delta(t))

A
  • the ½ is because 1 mol of H reacts with 1 mol of I2
    - ½ is related to stoichiometry of the reaction
    - notice with respect to the product there is no – coefficient
    - the rate of product is + as more product is created
    - if 100 I2 react per second then 200 HI molecules form per second
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9
Q

2:1 reaction rate

A

then for each 1 reactant(decreaing) 2 product(increasing)

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10
Q

the average rate of a reaction decreases as the reaction progresses

A

typically as the reactants transform to products, their concentrations decrease, and the reaction slows down

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11
Q

instantaneous rate

A

the rate of the reaction at any one point in time

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12
Q

instantaneous rate(at 50s)

A

= -(delta[H2])/delta(t) = -0.28M/40s = 0.0070 M/s

OR at 50s = +1/2(delta[HI]/delta(t)) = +1/2(.56M)/(40s) = 0.0070 M/s

  • the rate is the same with either one of the reactants or the product calculation
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13
Q

aA + bB -> cC + dD

A
  • A and B are reactants and C and D are products

- a,b,c,d are the stoichiometric coefficients

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14
Q

Rate =

A

-(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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15
Q

What does know the rate of change in the concentration of any one reactant or product at a point in time allow?

A

to determine the rate of change in the concentration of any other reactant or product at that point in time

  • predicting the rate of some future time is NOT possible from this equation
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16
Q

spectroscopy

A

as the intensity of the light absorption of a color decreases you can directly measure the concentration of a reactant as a function of time

  • can measure is femtoseconds(10^-15 second!)
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17
Q

changes in pressure(pressure measurement)

A

reactions in which the number of moles of a gaseous reactants and products changes as the reaction proceeds can be readily monitored

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18
Q

changes in pressure(pressure measurement):

2N2O(g) -> 2N2(g) + O2(g)

A
  • the pressure is going from 2 mols to 3 mols and this can be directly measured
    - the rise in pressure can be used to determine relative concentrations of reactants and products as a function of time.
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19
Q

polarity(polarimetry)

A

can be used to determine how the light rotates(clockwise or counterclockwise)
- the degree the light rotates can be measured

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20
Q

3 techniques to monitor a reaction as it occurs in a reaction vessel

A
  • polarimetry
  • spectroscopy
  • pressure measurement
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21
Q

rate law

A

Rate = k[A]^n

  • k = the constant of proportionality or the rate constant
  • n is the reaction order
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22
Q

rate law:

Zero order

A

Rate = k[A]^n

  • n = 0 then the reaction is zero order and the rate is independent of the concentration of A
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23
Q

rate law:

1st order

A

Rate = k[A]^n

  • n = 1 then the reaction is first order and the rate is directly proportional to the rate of A
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24
Q

rate law:

2nd order

A

Rate = k[A]^n

  • n = 2 then the reaction is second order and the rate is proportional to the square of the concentration of A
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25
Q

Zero order reaction

A
  • rate of reaction is independent of the concentration of the reactant
    - rate = k[A]^0 = k
    - the reactant decreases linearly with time
    - sublimation is zero order because only molecules at the surface can sublime so the concentration does not change
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26
Q

First order reaction

A
  • rate of the reaction is directly proportional to the concentration of the reactant
    - rate = k[A]^1
    - the rate slows down as reaction proceeds because the concentration of the reactant decreases
    - the rate is directly proportional to the concentration
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27
Q

Second order reaction

A
  • the rate of the reaction is proportional to the square of the concentration of the reactant
    - rate = k[A]^2
    - quadratic relationship
    - rate is proportional to the square of concentration
    - the order of a reaction can be determined only by experiment
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28
Q

method of initial rates

A
  • the initial rate, the rate for a short period of time at the beginning of the reaction, is measured by running the reaction several times with different initial reactant concentrations to determine the concentrations effect
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29
Q

overall order

A

rate = k[A]^m[B]^n

the sum of the exponents(m+n)

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30
Q

the rate law much always be determined by _____.

A

experiment

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31
Q

integrated rate law

A

the relationship between the concentrations of the reactants and time

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32
Q

first order integrated rate law

A
  • rate = -(delta [A]/delta t)
    - Rate = k[A]
    - -(delta [A]/delta t) = k[A]
    Ln[A]t = -kt + ln[A]0
    - y = mx + b and is linear
    Ln([A]t/[A]0) = -kt
    - [A]t = concentration A at any time t
    - k is the rate constant
    - [A]0 = initial concentration of A
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33
Q

Second Order Integrated Rate Law

A
  • Rate = k[A]^2
    - rate = -(delta [A]/delta t)
    - -(delta [A]/delta t) = k[A]^2
    (1/[A]t) = kt + (1/[A]0)
    - y = mx + b
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34
Q

Zero Ordered Integrated Rate Law

A
  • Rate = k[A]^0 = k
    - -(delta [A]/delta t) = k
    [A]t = -kt + [A]0
    - a straight line
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35
Q

half life(t1/2)

A

the required time for the concentration of a reactant to fall to one half of its initial value

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36
Q

First order reaction half life

A
  • t1/2 = (0.693)/t

- the time to halve the amount takes a constant amount of time

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37
Q

Second Order Reaction Half Life

A
  • t1/2 = 1/(k[A]0)
    - the half life depends on the initial concentration
    - as the amount decreases the half life becomes longer
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38
Q

Zero Order Reaction Half Life

A
  • t1/2 = [A]0/2k
    - half life depends on the initial concentration
    - the half life gets shorter as the amount decreases
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39
Q

reaction order and rate law must be determined

A

experimentally

40
Q
  • rate law relates the rate of the reaction to
A

the concentration of the reactants

41
Q
  • integrated rate law relates the concentration of the reactants to
A

time

42
Q

the half life is the time it takes for the concentration of a reactant to fall to

A

one half of its initial value

43
Q
  • the half life of a first order reaction is
A

independent of the initial concentration

44
Q
  • the half lives of zero order and second order reactions depend on
A

the initial concentration

45
Q

the rates of chemical reactions are highly sensitive to

A

temperature

46
Q

an increase in temperature results in an increase in k

A

rate = k[A]^n

47
Q

Arrhenius equation

A

k = A(e)^(-E/(RT))

		- E = activation energy
		- A = frequency factor
		- e = exponential factor
		- R = 8.314 J/mol*K
  • activation energy Ea – an energy barrier or hump that must be surmounted for reactants to be transformed into products
    - frequency factor(A) – the number of times that the reactants approach the activation barrier per unit time
48
Q

activation state or transition state

A

a molecule must go through a high energy intermediate state to go from reactant to product

49
Q

a reaction requires some energy put in before

A

it can go and become exothermic

50
Q

the higher the activation energy the

A

slower the reaction rate at a given temperature

51
Q

exponential factor

A

a number between 0 and 1 that represents the fraction of molecules that have enough energy to make it over the activation barrier on a given approach

	- 10^9/s would be 10^-7 at a certain temperature
		- 10^9/2 * 10^-7 = 10^2/s
52
Q

Exponential factor =

A

e^(-Ea/RT)

53
Q

Exponential factor:

A low activation energy and high temperature make the negative exponent small

A

the closer to 1 the easier it is for molecules to surpass the barrier

54
Q

frequency factor

A

the number of time that the reactants approach the activation barrier

55
Q

exponential factor

A

the fraction of the approaches that are successful in surmounting the activation barrier and forming products

56
Q

the exponential factor increases with increasing temperature but decreases with increasing activation energy

A
  • ln(AB) = ln A + ln B
    • ln e^x = x
    • ln k = -(Ea/R)(1/T) + ln A
      y = mx + b
57
Q

Arrhenius Plot

A

a plot of the natural log of the rate constant(ln k) versus the inverse of the temperature in kelvins(1/T) yields a straight line with a slope of –Ea/R and a y-intercept of ln A

58
Q

2 point Arrhenius plot

A

Ln (k2/k1) = Ea/R(1/T1 – 1/T2)

59
Q

collision model

A

a chemical reaction occurs after sufficiently energetic collision between two reactant molecules

this implies that the frequency factor should be the number of collisions per second but the frequency factors tend to be smaller

60
Q

the frequency factor can be split into two parts

A
  • k = (p)(z)e^(-Ea/RT)
    - p = orientation factor
    - z = collision frequency
61
Q

collision frequency(z)

A

the number of collisions that occur per unit time

- calculate for a gas phase reaction from the pressure of the gasses and the temperature of the reaction mixture

62
Q

orientation factor(p)

A

even with enough frequency and energy molecules must be properly aligned to cause a reaction so the proper bonds can break and form
- p = 0.16 means 16 out of every 100 sufficiently energetic collisions are actually forming products

63
Q

Harpoon mechanism

A

K can pass an electron to Br without actually colliding.

- p = 4.8 and the Br “reels” the electron in via columbic charges

64
Q

2 criteria for a chemical reaction to occur:

A

sufficient energy AND the correct orientation

65
Q

when we write out a chemical reaction we typically

A

write out the overall reaction not the individual steps

66
Q

H2(g) + 2ICl(g) -> 2HCl + I2

Really its this:

A
  • Step 1: H2(g) + ICl(g) _> HI(g) + HCl(g)

- Step 2: HI(g) + ICl(g) -> HCl(g) + I2(g)

67
Q

reaction mechanism

A

the series of initial chemical steps by which an overall chemical reaction occurs
- step 1 + step 2 + step 3 all written out

68
Q

elementary step

A

each individual step in a reaction mechanism

		- can not be broken down into simpler steps
		- literally occur as written
69
Q

reaction intermediate

A

a molecule that forms in one elementary step and is consumed in another
- may not be in the final product at all

70
Q

molecularity

A

the number of reactant particles involved in an elementary step

71
Q

most common types of molecularity

A
  • unimolecular - A-> products
  • bimolecular - A + B -> products
  • termolecular - A + B + C -> products
72
Q

termolecular

A

A + B + C -> products

elementary step in which 3 reactant particles collide and very rare due to very low probability of all 3 particles colliding at the same time

73
Q

rate law of elementary steps:

A + B -> products

A

Rate = k[A][B]

74
Q

rate law of elementary steps:

A + A -> products

A

Rate = k[A]^2

75
Q

rate law of elementary steps:

A -> products

A

Rate = k[A]

76
Q

rate law of elementary steps:

A + A + A

A

rate = k[A]^3

77
Q

rate law of elementary steps:

A + A + B

A

rate = k[A]^2[B]

78
Q

rate law of elementary steps:

A + B + C

A

rate = k[A][B][C]

79
Q

rate – determining step

A

the elementary step that is much slower than the other steps

	- a freeway with one lane open
	- limits the overall rate of the reaction and subsequently determines the rate law for the overall reaction
80
Q

to validate a proposed mechanism(mechanisms can not be proven!)

A
  • the elementary steps in the mechanism must sum to the overall reaction
    - the rate law predicted by the mechanism must be consistent with the experimentally observed rate law
81
Q
  • if the mechanism has a fast initial step then some other subsequent step in the mechanism will be the rate limiting step
A
  • the rate law predicted by the rate limiting step may contain reaction intermediates
82
Q
  • if the first step is fast the products may build up and begin reacting with each other until they reach an equilibrium indicated by a double arrow
A

reactants products

		- 2NO(g)  H2N2(g)   fast
		- H2(g) + N2O2(g) -> H2O(g) + N2O(g) slow(rate limiting)
		- N2O(g) + H2(g) -> N2(g) + H2O(g) fast
		- Overall:  2H2(g) + 2NO(g) -> 2H2O(g) + N2(g)
  • to be valid:
    - steps must sum to the overall reaction(true)
    - rate law predicted by the mechanism must be consistent with the experimentally observed law(fails this)
    - rate law contains an intermediate(N2O2) and the observed law did not
    - rate = k2[H2][N2O2]
    - experimentally observed = k[H2][NO]^2
    - because of the intermediate equilibrium in the first step we can express the concentration of the intermediate in terms of the reactants of the overall equation
    - k1[NO]^2 = k-1[N2O2]
    - [N2O2] = k1/k-1[NO]^2
    - rate = k2[H2][N2O2]
    = k2H2[NO]^2
    = (k2k1)/k-1[H2][NO]^2
    Rate = k[H2][NO]^2
83
Q

catalyst

A

a substance that increases the rate of a chemical reaction but is not consumed by the reaction

	- works by providing an alternative mechanism for the reaction
	- lowers the activation energy
84
Q

homogeneous catalysis

A

the catalyst exist in the same phase or state as the reactants

	- gas catalyst with gas phase reactants
		- catalytic destruction of ozone by Cl(g)
85
Q

heterogeneous catalysis

A

catalysis exists in a different phase than the reactants

- solid catalysis in a catalytic converting converts gas phase or solution phase reactants leaving vehicle

86
Q

hydrogenation

A

adding hydrogen to something

87
Q

Hydrogenation

4 steps:

A
  1. absorption
  2. diffusion
  3. reaction
  4. desorption
88
Q

Hydrogenation

4 steps:

  1. absorption
A

the reactants are absorbed onto the metal surface

89
Q

Hydrogenation

4 steps:

  1. diffusion
A

the reactants diffuse on the surface until they approach each other

90
Q

Hydrogenation

4 steps:

  1. reaction
A

the reactants react to form the products

91
Q

Hydrogenation

4 steps:

  1. desorption
A

the products desorb from the surface into the gas phase

92
Q

enzymes

A

biological catalysis that increase the rates of biochemical reactions

93
Q

enzymes:

active site

A

specific area on an enzyme where a substrate may attach

- highly specific to each specific substrate

94
Q

enzymes:

substrate

A

a reactant molecule that fits in a specific active site

95
Q

enzymes greatly lower the activation energy of the reaction

A
  • E + S ES fast

- ES -> E + P slow, rate limiting

Chem 2 Heider (9 decks)

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