principles of chemistry ii chem 1212 chapter 13 dr. augustine ofori agyeman assistant professor of...
TRANSCRIPT
![Page 1: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/1.jpg)
PRINCIPLES OF CHEMISTRY II
CHEM 1212
CHAPTER 13
DR. AUGUSTINE OFORI AGYEMANAssistant professor of chemistryDepartment of natural sciences
Clayton state university
![Page 2: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/2.jpg)
CHAPTER 13
CHEMICAL KINETICS
![Page 3: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/3.jpg)
- Chemical reactions occur when reactant species strike each other and interact to form products
Reaction kinetics is studied to
- improve production of materials- increase quality and quantity of products
- increase energy efficiency- minimize pollution
etc
RATES OF REACTIONS
![Page 4: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/4.jpg)
RATES OF REACTIONS
Rate = change per unit time
Rate of reaction = change in concentration per unit time
t
cRate
![Page 5: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/5.jpg)
For a chemical reaction
Reactant → Product
- Rate at which reactants are consumed or products are formed in a given period of time is given as
RATES OF REACTIONS
Δt
]Δ[reactant
Δt
Δ[product]Rate
Units: M/s
Square brackets represent molar concentrations [reactant] = reactant concentration[product] = product concentration
![Page 6: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/6.jpg)
Rate of appearance of product = rate of disappearance of reactant
- Reactant concentration decreases during reaction∆[reactant] is negative
- Product concentration increases during reaction∆[product] is positive
- Rate is always positive
- Rate can be measured by following the concentrations of reactants or products
RATES OF REACTIONS
![Page 7: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/7.jpg)
- Rate of reaction is generally not constant
- Rate of reaction changes over the course of reaction
- Concentration of reactants or products are measured at regulartime intervals
- A graph of concentration vs time may be plotted
- Instantaneous rate at a given time is the slope of the tangent to the curve at that time
- Average rate is measured rate over a time interval
INSTANTANEOUS AND AVERAGE RATES
![Page 8: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/8.jpg)
INSTANTANEOUS AND AVERAGE RATES
∆x
∆y
x
yslopeRateousInstantane
![Page 9: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/9.jpg)
- Rate depends on stoichiometry of the reaction
- Rate is the ratio of rate of change of a substance to its coefficient
Consider the reaction
2HBr(g) → H2(g) + Br2(g)
2 mol HBr : 1 mol of each product
REACTION STOICHIOMETRY
Δt
Δ[HBr]x
2
1
Δt
]Δ[H
Δt
]Δ[BrreactionofRate 22
![Page 10: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/10.jpg)
For the decomposition of HBr
2HBr(g) → H2(g) + Br2(g)
If HBr concentration is decreasing at a rate of 0.52 M/sWhat is the rate of the reaction?
What is the rate of appearance of H2 and Br2?
REACTION STOICHIOMETRY
M/s0.27M/s)0.52(x2
1
Δt
Δ[HBr]x
2
1reactionofRate
M/s0.27HBrmol2
BrorHmol1xHBr)M/s0.52(
Δt
]Δ[H
Δt
]Δ[BrreactionofRate 2222
![Page 11: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/11.jpg)
Factors Affecting Rate of Chemical Reaction
- Concentration of reactants
- Reaction temperature
- Physical nature of reactants
- Catalysts
RATES OF REACTIONS
![Page 12: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/12.jpg)
Concentration of Reactants
- An increase in the concentration of reactants causes an increase in the rate of reaction
- Collisions are more frequent in a given time for higher concentrations
RATES OF REACTIONS
![Page 13: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/13.jpg)
Reaction Temperature
- An increase in temperature of a system increases the average kinetic energy of the reacting molecules
- An increase in kinetic energy results in an increase in collisions in a given time
- The rate of a chemical reaction normally doubles for every 10 oC raise in temperature
RATES OF REACTIONS
![Page 14: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/14.jpg)
Physical State of Reactants: solid, liquid, or gas
solid-state reactants
liquid-state reactants
gaseous-state reactants
< <
Increasing rateof reaction
RATES OF REACTIONS
![Page 15: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/15.jpg)
Physical State of Reactants: solid, liquid, or gas
- Most frequent collisions occur in the gaseous state (the most freedom of movement of particles)
Solid-State Particle Size- Smaller particles have larger surface area and
higher reaction rates
- Extremely small particles may result in very fast reaction rates and may lead to explosion
RATES OF REACTIONS
![Page 16: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/16.jpg)
Catalysts
-Catalysts increase the rate of a reaction without being used up
RATES OF REACTIONS
![Page 17: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/17.jpg)
- Rate of reaction is strongly influenced by concentrationsof reacting species
- Rate is proportional to the product of the concentrations of the reactants each raised to some power
aA + bB → cC + dD
Rate = k[A]x[B]y
x and y are usually positive integersk = rate constant
RATE LAW
![Page 18: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/18.jpg)
aA + bB → cC + dD
Rate = k[A]x[B]y
- x and y are not necessarily coefficients of A and B- x and y are the orders of the reaction
- Described as xth order in A and yth order in B
If x = 1 and y = 2The reaction is first order in A and second order in B
Overall order = 1 + 2 = 3
RATE LAW
![Page 19: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/19.jpg)
Example
For the reaction
2NO2(g) + F2(g) → 2NO2F(g)
Rate = k[NO2][F2]
The reaction is first order in NO2 and first order in F2
RATE LAW
![Page 20: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/20.jpg)
The decomposition of nitrosyl chloride was studied:2NOCl(g) ↔ 2NO(g) + Cl2(g)The following data were obtained
INITIAL RATE OF REACTION
[NOCl]0 (molecules/cm3)
3.0 x 1016
2.0 x 1016
1.0 x 1016
4.0 x 1016
Initial Rate (molecules/cm3·s)
5.98 x 104
2.66 x 104
6.64 x 103
1.06 x 105
What is the rate law? Calculate the rate constant
Rate = k[NOCl]2, k = 6.64 x 10-29 cm3/molecules∙s
![Page 21: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/21.jpg)
The reaction below was studied at -10 oC2NO(g) + Cl2(g) → 2NOCl(g)The following data were obtained
INITIAL RATE OF REACTION
[NO]0 (mol/L)
0.100.100.20
Initial Rate (mol/L)
0.180.361.45
What is the rate law? Calculate the rate constant
Rate = k[NO]2[Cl2], k = 1.8 x 102 L2/mol2
[Cl2]0 (mol/L)
0.100.200.20
![Page 22: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/22.jpg)
- Rate of reaction decreases with time
- Rate of reaction eventually goes to zero
- Concentrations of reactants decrease
- Concentrations of products increase
CONCENTRATION AND TIME
![Page 23: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/23.jpg)
- Rates are independent of the concentrations of the reactants
R → product
Rate = k[R]0
Rate = k
- Called differential rate law
Unit of k = unit of reaction rate = M/s
ZERO-ORDER RATE LAW
![Page 24: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/24.jpg)
- Graph of concentration vs time
ZERO-ORDER RATE LAWC
once
ntr
atio
n
Time
![Page 25: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/25.jpg)
- Rates are independent of the concentrations of the reactants
R → product
[R]t = [R]0 - kt
- Called integrated rate law
Unit of k = unit of reaction rate = M/s
ExamplesMetabolism of ethyl alcohol in the body
Biochemical reactions involving enzymes
ZERO-ORDER RATE LAW
![Page 26: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/26.jpg)
- Graph of concentration vs time is a straight line
ZERO-ORDER RATE LAWC
once
ntr
atio
n
Time
Slope = −kIntercept = [R]0
![Page 27: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/27.jpg)
- A large value of k implies a fast reaction
- The half-life (t1/2) is also used to describe the speed of a reaction
- Half-life is the time needed for the concentration of a reactant to decrease to half its original value
- A short half-life indicates a fast reaction
HALF - LIFE
ZERO-ORDER RATE LAW
![Page 28: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/28.jpg)
At t = 0Initial concentration = [R]0
At half-life t = t1/2
[R]t = ½[R]0
- Substitute in zero-order equation and simplify
HALF - LIFE
ZERO-ORDER RATE LAW
![Page 29: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/29.jpg)
HALF - LIFE
- Using the zero-order rate equation
[R]t = [R]0 – kt
- Simplifying gives
ZERO-ORDER RATE LAW
2k
[R]t 0
1/2
- Half-life for zero-order depends on concentration
![Page 30: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/30.jpg)
ZERO-ORDER RATE LAW
The reaction A → B + C
is known to be zero order in A and to have a rate constant of 5.0 x 10-2 mol/L·s at 25 oC. An experiment was run at 25 oC
where [A]0 = 1.0 x 10-3 M.
a) What is the integrated rate law for this reaction?b) Calculate the half-life for the reaction.
c) Calculate the concentration of B after 5.0 x 10-3 s has elapsed.
a) [A] = [A]0 - ktb) 1.0 x 10-2 sc) 2.5 x 10-4 M
![Page 31: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/31.jpg)
- Rate is proportional to the concentration of the reactant
R → product
FIRST-ORDER RATE LAW
k[R]Δt
Δ[R]Rate
- Called the differential form of the rate law
- Relates differences in concentration and time
Unit of k = s-1
![Page 32: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/32.jpg)
- A graph of concentration vs time describes an exponential decay
FIRST-ORDER RATE LAWC
once
ntr
atio
n
Time
![Page 33: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/33.jpg)
- Rate is proportional to the concentration of the reactant
R → product
FIRST-ORDER RATE LAW
- Called the integrated form of the rate law (describes an exponential decay)
- Relates instantaneous concentrations
[R]t = concentration of R at any time[R]0 = initial concentration at t = 0
e = base of natural logarithms ≈ 2.718
kt0t e[R][R]
![Page 34: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/34.jpg)
FIRST-ORDER RATE LAW
From the first-order rate equation
ktln[R]ln[R] 0t
kt0t e[R][R]
Take natural logarithm on both sides and simplify
kt[R]
[R]ln
0
t
or
![Page 35: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/35.jpg)
FIRST-ORDER RATE LAWln
[Con
cen
trat
ion
]
Time
Slope = −kIntercept = ln[R]0
A graph of ln[R]t vs time is a straight line
![Page 36: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/36.jpg)
At t = 0Initial concentration = [R]0
At half-life t = t1/2
[R]t = ½[R]0
Substitute in first-order equation and simplify
HALF - LIFE
kt[R]
[R]ln
0
t kt[R]
[R]ln
0
t
FIRST-ORDER RATE LAW
![Page 37: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/37.jpg)
HALF - LIFE
kt[R]
[R]ln
0
t kt[R]
[R]ln
0
t
kt0t e[R][R]
k
.6930t1/2
From the first-order rate equation
Substitute and simplify
FIRST-ORDER RATE LAW
![Page 38: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/38.jpg)
- Half-life of a first-order reaction is independent of the concentration of the reactant
- Depends on only the rate constant (k)
- Constant half-life from concentration vs time plot indicates first-order reaction
ExampleRadioactive decay processes
HALF - LIFE
kt[R]
[R]ln
0
t kt[R]
[R]ln
0
t
FIRST-ORDER RATE LAW
![Page 39: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/39.jpg)
The radioactive isotope 32P decays by first-order kinetics and has a half-life of 14.3 days. How long does it take for 95% of
a sample of 32P to decay?
k = 0.0485 1/dayt = 61.8 days
FIRST-ORDER RATE LAW
kt[R]
[R]ln
0
t kt[R]
[R]ln
0
t
![Page 40: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/40.jpg)
A first-order reaction is 75.0% complete in 320 second.a) What are the first and second half-lives for this reaction?
b) How long does it take for 90% completion?
a) 160 s for both first and second half-livesb) 532 s
FIRST-ORDER RATE LAW
kt[R]
[R]ln
0
t kt[R]
[R]ln
0
t
![Page 41: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/41.jpg)
Calculate the half-life of a first order reaction if the concentration of the reactant is 0.0451 M at 30.5 seconds after
the reaction starts and is 0.0321 M at 45.0 seconds after the reaction starts. How many seconds after the start of the
reaction does it take for the reactant concentration to decrease to 0.0100 M?
a) 29.5 sb) 94.9 s
kt[R]
[R]ln
0
t kt[R]
[R]ln
0
t
FIRST-ORDER RATE LAW
![Page 42: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/42.jpg)
- Rate is proportional to the concentration of the reactant raised to the second power
R → product
SECOND-ORDER RATE LAW
- Called the differential form of the rate law
- Relates differences in concentration and time
Unit of k = M-1s-1 or L/mol·s
2k[R]Δt
Δ[R]Rate
![Page 43: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/43.jpg)
- Graph of concentration vs time
SECOND-ORDER RATE LAWC
once
ntr
atio
n
Time
![Page 44: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/44.jpg)
- Rate is proportional to the concentration of the reactant raised to the second power
R → product
SECOND-ORDER RATE LAW
- Called the integrated form of the rate law
Unit of k = M-1s-1 or L/mol·s
kt[R]
1
[R]
1
0t
![Page 45: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/45.jpg)
- A graph of 1/concentration vs time is a straight line
SECOND-ORDER RATE LAW1/
[Con
cen
trat
ion
]
Time
Slope = kIntercept = 1/[R]0
![Page 46: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/46.jpg)
HALF-LIFE
Half-life depends on starting concentration
01/2 k[R]
1t
SECOND-ORDER RATE LAW
![Page 47: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/47.jpg)
For the reaction A → products
successive half-lives are observed to be 10.0, 20.0, and 40.0 min for an experiment in which [A]0 = 0.10 M. Calculate the
concentration of A at a) 30.0 minb) 70.0 minc) 80.0 min
a) 0.025 Mb) 0.013 Mc) 0.011 M
kt[R]
[R]ln
0
t kt[R]
[R]ln
0
t
SECOND-ORDER RATE LAW
![Page 48: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/48.jpg)
SECOND-ORDER RATE LAW
Consider the following initial rate data for the decomposition of compound AB to give A and B
[AB]0, mol/L: 0.200 0.400 0.600Initial rate, mol/L·s: 3.20 x 10-3 1.28 x 10-2 2.88 x 10-2
Determine the half-life for the decomposition reaction initially having 1.00 M AB present
Rate = k[AB]2
k = 0.0800 L/mol·st1/2 = 12.5 s
![Page 49: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/49.jpg)
RATE AND TEMPERATURE
- Almost all reactions go faster at higher temperatures
- The rate of most reactions increase at increasing temperature
- The order of the reaction usually does not change with temperature
![Page 50: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/50.jpg)
RATE AND TEMPERATURE
Example
For the reactionNO(g) + O3(g) → NO2(g) + O2(g)
The rate constant increases with increasing temperature
k (L
/mol
·s)
T (K)
![Page 51: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/51.jpg)
COLLISION THEORY
- Explains the rate of reactions in terms of molecular-scale collisions
- The basic assumption is that molecules must collide to react
- The collision frequency (Z) is the number of collisions per second
- Z depends on the concentrations of the reacting species
![Page 52: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/52.jpg)
COLLISION THEORY
- The collision frequency (Z) between two molecules is proportional to the product of their concentrations
- For two reacting molecules XY and AB
Z α [XY][AB]
Z = Z0[XY][AB]
Z0 = is the proportionality constant
Z0 depends on sizes and speed of reacting species
![Page 53: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/53.jpg)
COLLISION THEORY
- Collision frequency increases with increasing temperature as molecules move faster
- However, the increase in collision frequency cannot accountfor the temperature dependence of reaction rate
- Not every collision results in a chemical reaction
![Page 54: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/54.jpg)
ACTIVATION ENERGY (Ea)
- Not all collisions result in the formation of products (by Svante Arrhenius)
- Molecules must collide with enough energy to rearrange the bonds
- Molecules bounce off if the total energy of colliding species is not enough
- Activation energy (Ea) is the minimum collision energy required for a reaction to occur
![Page 55: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/55.jpg)
THE ACTIVATED COMPLEX
- Is a transition state
- The least stable or highest energy transition state
- Very unstable and concentration is extremely small
- The energy needed to from the activated complex from the reactants is the activation energy
- Reactions with high activation energies are generally slower than reactions with low activation energies
![Page 56: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/56.jpg)
Reaction coordinate
pote
nti
al e
nerg
y
ENERGY LEVEL DIAGRAM
Products
Activated complex
ReactantsEa
![Page 57: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/57.jpg)
EFFECT OF TEMPERATURE
- The effect of temperature on reaction rate is influenced by the magnitude of the activation energy
- The number of molecules with high enough kinetic energies to initiate a reaction is directly related to temperature
- The fraction of collisions (fr) with energy in excess of Ea
/RTEr
aef
R is the gas constant = 8.314 J/mol·KT is the temperature in Kelvin
![Page 58: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/58.jpg)
EFFECT OF TEMPERATURE
- fr is between 0 and 1
- fr gets closer to 1 as T increases
- Ea does not change with T
- Number of collisions exceeding Ea increases exponentially with T
![Page 59: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/59.jpg)
Energy
Fra
ctio
n
High Temperature Gas
Low Temperature Gas
ENERGY DISTRIBUTION IN GAS MOLECULES
Ea
![Page 60: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/60.jpg)
EFFECT OF TEMPERATURE
Rate of reaction = (collision frequency) x (fraction exceeding Ea)
Rate = Z x fr
Z = Z0[XY][AB]
Experimental rate = k[XY][AB] /RTE
oa[XY][AB]eZratePredicted
/RTEo
aeZk
k is the rate constant
![Page 61: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/61.jpg)
STERIC FACTOR
- The expression predicts faster rates than experimentally observed
- Not all collisions with energies greater than Ea result in a reaction
- The correct orientation of reactants is an important factor
- The steric factor (p) expresses the need for the correct orientation
Rate = (steric factor) x (collision frequency) x (fraction exceeding Ea)
/RTEo
a[XY][AB]eZRate p
![Page 62: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/62.jpg)
STERIC FACTOR
A = pZo
A is known as the pre-exponential term
/RTEa[XY][AB]eRate A
/RTEaek A
The Arrhenius equation
- A includes the steric factor and cannot be predicted by theory- A can only be determined by experiment
![Page 63: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/63.jpg)
THE ARRHENIUS EQUATION/RTEaek A
Take natural log of both sides
/RTEAlnkln a
For an Arrhenius plot
- That is a graph of ln k versus 1/T
Slope = -Ea/R
Intercept = ln A
![Page 64: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/64.jpg)
THE ARRHENIUS EQUATION
/RTEaek A
Consider rate constants k1 and k2 at temperatures T1 and T2
21
a
2
1
T
1
T
1
R
E
k
kln
![Page 65: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/65.jpg)
THE ARRHENIUS EQUATION
The activation energy for the decomposition of HI(g) to H2(g) and I2(g) is 186 kJ/mol. The rate constant at 555 K is
3.52 x 10-7 L/mol·s. What is the rate constant at 645 K?
9.60 x 10-5 L/mol·s
![Page 66: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/66.jpg)
THE ARRHENIUS EQUATION
A first order reaction has rate constant of 4.6 x 10-2 s-1 and 8.1 x 10-2 s-1 at 0 oC and 20 oC, respectively. What is the value of the activation energy?
Ea = 19 kJ/mol
![Page 67: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/67.jpg)
THE ARRHENIUS EQUATION
A certain reaction has an activation energy of 54.0 kJ/mol.As the temperature is increased from 22 oC to a higher temperature,
the rate constant increases by a factor of 7.00. Calculate the higher temperature.
T2 = 324 K or 51 oC
![Page 68: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/68.jpg)
CATALYSIS
Rate of reaction can be increased in two ways
1) Increase the temperature
2) Reduce the activation energy or increase the steric factor(addition of catalyst)
- A catalyst is a substance that increases the rate of reaction but is not consumed in the reaction
- A catalyzed reaction generally has lower activation energy
![Page 69: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/69.jpg)
- Catalysts increase the rate of a reaction without being used up
- Provide alternative reaction pathways with lower activation energies
Uncatalyzed reaction: X + Y → XY
Catalyzed reaction: Step 1 X + C → XCStep 2 XC + Y → XY + C
CATALYSIS
![Page 70: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/70.jpg)
Reaction pathway
pote
nti
al e
nerg
y uncatalyzed activation
energy
catalyzed activation
energy
CATALYSIS
![Page 71: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/71.jpg)
HOMOGENEOUS CATALYSIS
- Present in the same phase as the reactants
ExampleN2(g) + O2(g) → 2NO(g)
The formation of ozone
![Page 72: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/72.jpg)
HETEROGENEOUS CATALYSIS
- Present in a different phase from the reactants
ExamplesUse of solid metal catalysts such as platinum, nickel, palladium, titanium
Use of platinum catalyst for the production of methanol from hydrogen and carbon monoxide
2H2(g) + CO(g) → CH3OH(g)
- Catalysts can determine the nature of products formedPlatinum catalyst produces methanol
Nickel catalyst produces methane and water
![Page 73: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/73.jpg)
ENZYME CATALYSIS
- Enzymes are large molecules that catalyze specific biochemical reactions
- An enzyme is specifically tailored to facilitate a given reaction
- Enzymes increase the rate of reaction by increasing the steric factor rather than decreasing the activation energy
- Enzymes are generally named after the reactions they catalyze(that is their functions)
ExamplesCarboxypeptidase-A, Alcohol dehydrogenase (ADH)
![Page 74: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/74.jpg)
COLLISIONS BETWEEN MOLECULES
- The sequence of steps leading from reactants to products is known as the reaction mechanism
- Some reactions require only one step (a single collision)
- Other reactions require more than one collisions leading to the formation of intermediates
![Page 75: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/75.jpg)
INTERMEDIATES
- Compounds that are produced in one step and consumed in another
- Not observed among the products of the reaction
- Differ from activated complex
- An intermediate is in a shallow minimum in the energy level diagram
- An activated complex occurs at the maximum in the energy level diagram
![Page 76: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/76.jpg)
Reaction coordinate
pote
nti
al e
nerg
yENERGY LEVEL DIAGRAM
Products
Intermediates
Reactants
![Page 77: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/77.jpg)
ELEMENTARY STEP
- Chemical equation that describes an actual molecular-level event
- The overall reaction is the sum of the elementary reactions
ExampleNO2 + NO2 → NO3 + NO step 1NO3 + CO → NO2 + CO2 step 2
NO2 + CO → NO + CO2 overall
NO3 is an intermediate
![Page 78: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/78.jpg)
RATE LAW FOR ELEMENTARY STEP
- Rate law of an elementary step can be written directly from the stoichiometry of that step
Consider an elementary stepiA + jB → products
Rate = k[A]i[B]j
- The rate law for an overall reaction cannot be determined from the stoichiometry
![Page 79: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/79.jpg)
Molecularity- The number of species involved in a single elementary step
Unimolecular Step- Involves the spontaneous decomposition of a single molecule
- First-order rate law describes the kineticsHCl → H + Cl
Bimolecular Step- Involves the collision of two species
- Second-order rate law describes the kineticsNO2 + NO2 → N2O4
RATE LAW FOR ELEMENTARY STEP
![Page 80: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/80.jpg)
Termolecular Step- Involves the collision of three species
- Third-order rate law describes the kinetics- Uncommon
NO2 + NO + O2 → NO3 + NO2
- Collisions involving four or more species are very rare
RATE LAW FOR ELEMENTARY STEP
![Page 81: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/81.jpg)
RATE-LIMITING STEP
- The slowest elementary step in a given reaction
- The rate of a chemical reaction is limited by the rate of the slowest step
- The rate law of the slowest step is consistent with the experimental rate law of the overall reaction
Example2NO ↔ N2O2 fast, reversible
N2O2 + Cl2 → 2NOCl slow step (rate limiting)
![Page 82: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/82.jpg)
COMPLEX REACTION MECHANISMS
- Reactions in which the rate limiting step is not the first step
- Reaction rate may depend on intermediates
- Intermediates are unstable and their concentrations are difficult to measure
- Rate laws are not written in terms of intermediates
- Other complex reactions contain rapid and reversible steps before the rate-limiting step
![Page 83: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 13 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state](https://reader035.vdocuments.us/reader035/viewer/2022062803/56649caf5503460f94973644/html5/thumbnails/83.jpg)
ENZYME METABOLISM
- Many enzyme catalyzed reactions follow the Michaelis-Menten mechanism
E + S ↔ ES → E + P
- Rate of reaction is zero order in substrate (S)
Substrate- The compound on which the enzyme acts
- Product (P) does not bind to enzyme (E)- First step is fast and reversible
- Second step is irreversible