Ch 13 - Reactions Flashcards
ectotherms
organisms with their body temperature dependent on their surroundings
What is the effect of cold on reptiles?
the rate of reactions in their muscles occur more slowly making it impossible to move quickly
How does the rate of a reaction affect a rocket?
extremely important.
Too slow and no lift.
Too fast and can explode.
What does the rate of a chemical reaction measure?
How fast the reaction occurs
A fast reaction rate =
a larger fraction of the molecules react to form products in a given period of time.
H2(g) + I2(g) -> 2HI(g)
- 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
Why is the reaction rate defined as negative?
- 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)
Rate with respect to the product = +(1/2)(delta[HI]/delta(t))
- 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
2:1 reaction rate
then for each 1 reactant(decreaing) 2 product(increasing)
the average rate of a reaction decreases as the reaction progresses
typically as the reactants transform to products, their concentrations decrease, and the reaction slows down
instantaneous rate
the rate of the reaction at any one point in time
instantaneous rate(at 50s)
= -(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
aA + bB -> cC + dD
- A and B are reactants and C and D are products
- a,b,c,d are the stoichiometric coefficients
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))
What does know the rate of change in the concentration of any one reactant or product at a point in time allow?
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
spectroscopy
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!)
changes in pressure(pressure measurement)
reactions in which the number of moles of a gaseous reactants and products changes as the reaction proceeds can be readily monitored
changes in pressure(pressure measurement):
2N2O(g) -> 2N2(g) + O2(g)
- 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.
polarity(polarimetry)
can be used to determine how the light rotates(clockwise or counterclockwise)
- the degree the light rotates can be measured
3 techniques to monitor a reaction as it occurs in a reaction vessel
- polarimetry
- spectroscopy
- pressure measurement
rate law
Rate = k[A]^n
- k = the constant of proportionality or the rate constant
- n is the reaction order
rate law:
Zero order
Rate = k[A]^n
- n = 0 then the reaction is zero order and the rate is independent of the concentration of A
rate law:
1st order
Rate = k[A]^n
- n = 1 then the reaction is first order and the rate is directly proportional to the rate of A
rate law:
2nd order
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
Zero order reaction
- 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
First order reaction
- 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
Second order reaction
- 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
method of initial rates
- 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
overall order
rate = k[A]^m[B]^n
the sum of the exponents(m+n)
the rate law much always be determined by _____.
experiment
integrated rate law
the relationship between the concentrations of the reactants and time
first order integrated rate law
- 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
Second Order Integrated Rate Law
- 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
Zero Ordered Integrated Rate Law
- Rate = k[A]^0 = k
- -(delta [A]/delta t) = k
[A]t = -kt + [A]0
- a straight line
half life(t1/2)
the required time for the concentration of a reactant to fall to one half of its initial value
First order reaction half life
- t1/2 = (0.693)/t
- the time to halve the amount takes a constant amount of time
Second Order Reaction Half Life
- t1/2 = 1/(k[A]0)
- the half life depends on the initial concentration
- as the amount decreases the half life becomes longer
Zero Order Reaction Half Life
- t1/2 = [A]0/2k
- half life depends on the initial concentration
- the half life gets shorter as the amount decreases
reaction order and rate law must be determined
experimentally
- rate law relates the rate of the reaction to
the concentration of the reactants
- integrated rate law relates the concentration of the reactants to
time
the half life is the time it takes for the concentration of a reactant to fall to
one half of its initial value
- the half life of a first order reaction is
independent of the initial concentration
- the half lives of zero order and second order reactions depend on
the initial concentration
the rates of chemical reactions are highly sensitive to
temperature
an increase in temperature results in an increase in k
rate = k[A]^n
Arrhenius equation
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
activation state or transition state
a molecule must go through a high energy intermediate state to go from reactant to product
a reaction requires some energy put in before
it can go and become exothermic
the higher the activation energy the
slower the reaction rate at a given temperature
exponential factor
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
Exponential factor =
e^(-Ea/RT)
Exponential factor:
A low activation energy and high temperature make the negative exponent small
the closer to 1 the easier it is for molecules to surpass the barrier
frequency factor
the number of time that the reactants approach the activation barrier
exponential factor
the fraction of the approaches that are successful in surmounting the activation barrier and forming products
the exponential factor increases with increasing temperature but decreases with increasing activation energy
- ln(AB) = ln A + ln B
- ln e^x = x
- ln k = -(Ea/R)(1/T) + ln A
y = mx + b
Arrhenius Plot
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
2 point Arrhenius plot
Ln (k2/k1) = Ea/R(1/T1 – 1/T2)
collision model
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
the frequency factor can be split into two parts
- k = (p)(z)e^(-Ea/RT)
- p = orientation factor
- z = collision frequency
collision frequency(z)
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
orientation factor(p)
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
Harpoon mechanism
K can pass an electron to Br without actually colliding.
- p = 4.8 and the Br “reels” the electron in via columbic charges
2 criteria for a chemical reaction to occur:
sufficient energy AND the correct orientation
when we write out a chemical reaction we typically
write out the overall reaction not the individual steps
H2(g) + 2ICl(g) -> 2HCl + I2
Really its this:
- Step 1: H2(g) + ICl(g) _> HI(g) + HCl(g)
- Step 2: HI(g) + ICl(g) -> HCl(g) + I2(g)
reaction mechanism
the series of initial chemical steps by which an overall chemical reaction occurs
- step 1 + step 2 + step 3 all written out
elementary step
each individual step in a reaction mechanism
- can not be broken down into simpler steps - literally occur as written
reaction intermediate
a molecule that forms in one elementary step and is consumed in another
- may not be in the final product at all
molecularity
the number of reactant particles involved in an elementary step
most common types of molecularity
- unimolecular - A-> products
- bimolecular - A + B -> products
- termolecular - A + B + C -> products
termolecular
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
rate law of elementary steps:
A + B -> products
Rate = k[A][B]
rate law of elementary steps:
A + A -> products
Rate = k[A]^2
rate law of elementary steps:
A -> products
Rate = k[A]
rate law of elementary steps:
A + A + A
rate = k[A]^3
rate law of elementary steps:
A + A + B
rate = k[A]^2[B]
rate law of elementary steps:
A + B + C
rate = k[A][B][C]
rate – determining step
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
to validate a proposed mechanism(mechanisms can not be proven!)
- 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
- if the mechanism has a fast initial step then some other subsequent step in the mechanism will be the rate limiting step
- the rate law predicted by the rate limiting step may contain reaction intermediates
- 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
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
catalyst
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
homogeneous catalysis
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)
heterogeneous catalysis
catalysis exists in a different phase than the reactants
- solid catalysis in a catalytic converting converts gas phase or solution phase reactants leaving vehicle
hydrogenation
adding hydrogen to something
Hydrogenation
4 steps:
- absorption
- diffusion
- reaction
- desorption
Hydrogenation
4 steps:
- absorption
the reactants are absorbed onto the metal surface
Hydrogenation
4 steps:
- diffusion
the reactants diffuse on the surface until they approach each other
Hydrogenation
4 steps:
- reaction
the reactants react to form the products
Hydrogenation
4 steps:
- desorption
the products desorb from the surface into the gas phase
enzymes
biological catalysis that increase the rates of biochemical reactions
enzymes:
active site
specific area on an enzyme where a substrate may attach
- highly specific to each specific substrate
enzymes:
substrate
a reactant molecule that fits in a specific active site
enzymes greatly lower the activation energy of the reaction
- E + S ES fast
- ES -> E + P slow, rate limiting
