Upcoming SlideShare
×

Thanks for flagging this SlideShare!

Oops! An error has occurred.

×
Saving this for later? Get the SlideShare app to save on your phone or tablet. Read anywhere, anytime – even offline.
Standard text messaging rates apply

# General chemistry ii chapter 14

2,067

Published on

Published in: Education, Technology
1 Like
Statistics
Notes
• Full Name
Comment goes here.

Are you sure you want to Yes No
• Be the first to comment

Views
Total Views
2,067
On Slideshare
0
From Embeds
0
Number of Embeds
0
Actions
Shares
0
121
0
Likes
1
Embeds 0
No embeds

No notes for slide

### Transcript

• 1. Chapter 14 Chemical Kinetics CHEMISTRY The Central Science 9th Edition
• 2.
• Kinetics is the study of how fast chemical reactions occur.
• There are 4 important factors which affect the rates of chemical reactions:
• reactant concentration,
• temperature,
• action of catalysts, and
• surface area.
Kinetics
• 3.
• The speed of a reaction is defined as the change that occurs per unit time.
• It is determined by measuring the change in concentration of a reactant or product with time.
• The speed the of the reaction is called the reaction rate .
• For a reaction A  B
• Suppose A reacts to form B. Let us begin with 1.00 mol A.
Reaction Rates
• 4. Change in Concentration of Reactions 10 20
• 5.
• At t = 0 (time zero) there is 1.00 mol A (100 red spheres) and no B present.
• At t = 10 min, there is 0.54 mol A and 0.26 mol B.
• At t = 20 min, there is 0.30 mol A and 0.70 mol B.
• Calculating,
Calculating Reaction Rates Using Products
• 6.
• For the reaction A  B there are two ways of measuring rate:
• the speed at which the products appear (i.e. change in moles of B per unit time), or
• the speed at which the reactants disappear (i.e. the change in moles of A per unit time).
• The equation, when calculating rates of reactants, is multiplied by -1 to compensate for the negative concentration.
• By convention rates are expressed as positive numbers.
Calculating Reaction Rates Using Reactants
• 7.
• Most useful units for rates are to look at molarity. Since volume is constant, molarity and moles are directly proportional.
• Consider:
• C 4 H 9 Cl( aq ) + H 2 O( l )  C 4 H 9 OH( aq ) + HCl( aq )
Reaction Rates Reactants Products
• 8. Reaction Rates for C 4 H 9 Cl
• 9.
• C 4 H 9 Cl( aq ) + H 2 O( l )  C 4 H 9 OH( aq ) + HCl( aq )
• We can calculate the average rate in terms of the disappearance of C 4 H 9 Cl.
• The units for average rate are mol/L·s or M/s .
• The average rate decreases with time.
• We plot [C 4 H 9 Cl] versus time.
• The rate at any instant in time (instantaneous rate) is the slope of the tangent to the curve.
• Instantaneous rate is different from average rate.
• We usually call the instantaneous rate the rate.
Properties of C 4 H 9 Cl Reaction
• 10. Instantaneous Reaction Rates for C 4 H 9 Cl
• 11.
• For the reaction
• C 4 H 9 Cl( aq ) + H 2 O( l )  C 4 H 9 OH( aq ) + HCl( aq )
• we know
• In general for
• a A + b B  c C + d D
Reaction Rates and Stoichiometry
• 12.
• In general rates increase as concentrations increase.
• NH 4 + ( aq ) + NO 2 - ( aq )  N 2 ( g ) + 2H 2 O( l )
Concentration and Rate Table
• 13.
• For the reaction
• NH 4 + ( aq ) + NO 2 - ( aq )  N 2 ( g ) + 2H 2 O( l )
• we note
• as [NH 4 + ] doubles with [NO 2 - ] constant the rate doubles,
• as [NO 2 - ] doubles with [NH 4 + ] constant, the rate doubles,
• We conclude rate  [NH 4 + ][NO 2 - ].
• Rate law:
• The constant k is the rate constant.
Concentration and Rate Equation
• 14.
• For a general reaction with rate law
• we say the reaction is m th order in reactant 1 and n th order in reactant 2.
• The overall order of reaction is m + n + ….
• A reaction can be zeroth order if m , n , … are zero.
• Note the values of the exponents (orders) have to be determined experimentally. They are not simply related to stoichiometry.
Exponents in the Rate Law
• 15.
• A reaction is zero order in a reactant if the change in concentration of that reactant produces no effect.
• A reaction is first order if doubling the concentration causes the rate to double.
• A reaction is n th order if doubling the concentration causes an 2 n increase in rate.
• Note that the rate constant does not depend on concentration.
Determining Order of Reactions
• 16.
• First Order Reactions
• Goal: convert rate law into a convenient equation to give concentrations as a function of time.
• For a first order reaction, the rate doubles as the concentration of a reactant doubles.
Concentration Change with Time
• 17.
• A plot of ln[A] t versus t is a straight line with slope - k and intercept ln[A] 0 .
• Put into simple math language : y = m x + b
• Plotting of this equation for given reaction should yield a straight line if the reaction is first order.
• In the above we use the natural logarithm, ln, which is log to the base e .
Plotting First Order Reactions
• 18.
• Left graph plotted with pressure vs. t, and right graph plotted with ln(pressure) vs. t.
Example Plots of a 1 st Order Reaction
• 19.
• Second Order Reactions
• For a second order reaction with just one reactant
• A plot of 1/[A] t versus t is a straight line with slope k and intercept 1/[A] 0
• Put into simple math language: y = m x + b
• For a second order reaction, a plot of ln[A] t vs. t is not linear.
• However, a plot of 1/[A] t versus t is a straight line
Concentration Change with Time
• 20.
• Left is Ln[NO 2 ] vs t, and right is 1/[NO 2 ] vs. t.
Example plots of a 2 nd Order Reaction
• 21.
• First Order Reactions
• Half-life is the time taken for the concentration of a reactant to drop to half its original value.
• For a first order process, half life, t ½ is the time taken for [A] 0 to reach ½[A] 0 .
• Mathematically defined by:
• The half-life for a 1 st order reactions depends only on k
Half-Life Reactions
• 22.
• Second Order
• Mathematically defined by:
• A second order reaction’s half-life depends on the initial concentration of the reactants
Half-Life Reaction
• 23.
• The Collision Model
• Most reactions speed up as temperature increases. (E.g. food spoils when not refrigerated.)
• When two light sticks are placed in water: one at room temperature and one in ice, the one at room temperature is brighter than the one in ice.
• The chemical reaction responsible for chemiluminescence is dependent on temperature: the higher the temperature, the faster the reaction and the brighter the light.
Temperature and Rate
• 24. The Collision Model
• As temperature increases, the rate increases.
• 25.
• Goal: develop a model that explains why rates of reactions increase as concentration and temperature increases.
• The collision model: in order for molecules to react they must collide.
• The greater the number of collisions the faster the rate.
• The more molecules present, the greater the probability of collision and the faster the rate.
• Faster moving molecule collide with greater energy and more frequently, increasing reaction rates.
Collision Model: The Central Idea
• 26.
• The Collision Model
• The higher the temperature, the more energy available to the molecules and the faster the rate.
• Complication: not all collisions lead to products. In fact, only a small fraction of collisions lead to product.
• Why is this?
• The Orientation Factor
• In order for reaction to occur the reactant molecules must collide in the correct orientation and with enough energy to form products.
The Speed of a Reaction
• 27.
• Consider:
• Cl + NOCl  NO + Cl 2
• There are two possible ways that Cl atoms and NOCl molecules can collide; one is effective and one is not.
The Orientation Factor
• 28.
• Arrhenius: molecules must posses a minimum amount of energy to react. Why?
• In order to form products, bonds must be broken in the reactants.
• Bond breakage requires energy.
• Activation energy, E a , is the minimum energy required to initiate a chemical reaction.
Activation Energy
• 29. Energy Profile for Methly Isonitrile
• 30.
• How does a methyl isonitrile molecule gain enough energy to overcome the activation energy barrier?
• From kinetic molecular theory, we know that as temperature increases, the total kinetic energy increases.
• We can show the fraction of molecules, f , with energy equal to or greater than E a is
• where R is the gas constant (8.314 J/mol·K).
Fraction of Molecules Possessing E a
• 31. Activation Energy, E a , Plot
• 32.
• Arrhenius discovered most reaction-rate data obeyed the Arrhenius equation:
• k is the rate constant, E a is the activation energy, R is the gas constant (8.314 J/K-mol) and T is the temperature in K.
• A is called the frequency factor.
• A is a measure of the probability of a favorable collision.
• Both A and E a are specific to a given reaction.
The Arrhenius Equation
• 33.
• If we have a lot of data, we can determine E a and A graphically by rearranging the Arrhenius equation:
• From the above equation, a plot of ln k versus 1/ T will have slope of – E a /R and intercept of ln A .
Determing Activation Energy
• 34. Ln k versus 1/T
• E a can be determined by finding the slope of the line.
• 35.
• The balanced chemical equation provides information about the beginning and end of reaction.
• The reaction mechanism gives the path of the reaction (i.e., process by which a reaction occurs).
• Mechanisms provide a very detailed picture of which bonds are broken and formed during the course of a reaction.
• Elementary Steps
• Elementary step: any process that occurs in a single step.
Reaction Mechanisms
• 36.
• Molecularity: the number of molecules present in an elementary step.
• Unimolecular: one molecule in the elementary step,
• Bimolecular: two molecules in the elementary step, and
• Termolecular: three molecules in the elementary step.
• It is not common to see termolecular processes (statistically improbable).
Properties of the Elementary Step Process
• 37.
• Some reaction proceed through more than one step:
• Consider the reaction of NO 2 and CO
• NO 2 ( g ) + NO 2 ( g )  NO 3 ( g ) + NO( g )
• NO 3 ( g ) + CO( g )  NO 2 ( g ) + CO 2 ( g )
• Notice that if we add the above steps, we get the overall reaction:
• NO 2 ( g ) + CO( g )  NO( g ) + CO 2 ( g )
• The elementary steps must add to give the balanced chemical equation.
• Intermediate : a species which appears in an elementary step which is not a reactant or product
Multistep Mechanisms
• 38.
• Rate Laws for Elementary Steps
• The rate law of an elementary step is determined by its molecularity:
• Unimolecular processes are first order,
• Bimolecular processes are second order, and
• Termolecular processes are third order.
• Rate Laws for Multistep Mechanisms
• Rate-determining step: is the slowest of the elementary steps.
Reaction Mechanisms
• 39.
• Rate Laws for Elementary Steps
Table 14.3, Page 551
• 40.
• It is possible for an intermediate to be a reactant.
• Consider
• 2NO( g ) + Br2( g )  2NOBr( g )
• The experimentally determined rate law is
• Rate = k[NO] 2 [Br2]
• Consider the following mechanism
Initial Fast Step of a Mechanisms
• 41.
• The rate law is (based on Step 2 ):
• Rate = k 2 [NOBr 2 ][NO]
• The rate law should not depend on the concentration of an intermediate (intermediates are usually unstable).
• Assume NOBr 2 is unstable, so we express the concentration of NOBr 2 in terms of NOBr and Br 2 assuming there is an equilibrium in step 1 we have
Intermediate as a Reactant
• 42.
• By definition of equilibrium:
• Therefore, the overall rate law becomes
• Note the final rate law is consistent with the experimentally observed rate law.
Intermediate as a Reactant Conts.
• 43.
• A catalyst changes the rate of a chemical reaction.
• Catalyst lower the overall E a for a chemical reaction.
• There are two types of catalyst:
• Homogeneous and heterogeneous
• Example: Cl atoms are catalysts for the destruction of ozone.
• Homogeneous Catalysis
• The catalyst and reaction is in one phase.
• Heterogeneous Catalysis
• The catalyst and reaction exists in a different phase.
Catalysis
• 44. The Effects of a Catalyst
• 45.
• Catalysts can operate by increasing the number of effective collisions (i.e., from the Arrhenius equation: catalysts increase k which results in increasing A or decreasing E a .
• A catalyst may add intermediates to the reaction.
• Example: In the presence of Br - , Br 2 ( aq ) is generated as an intermediate in the decomposition of H 2 O 2 .
Functions of the Catalysis
• 46. End of Chapter 14 Chemical Kinetics