kinetics review Flashcards
Kinetics Rates of RXN intro
Thus far, we have been discussing whether a reaction is spontaneous or not. The spontaneity of a reaction does not predict how fast a reaction will occur. The rate of reaction determines how fast or slow a reaction will occur.
Kinetics: Rates of Reaction
Why change the rate of a reaction?
- To speed up or slow down a reaction
- Slow the rate of rusting
- Slow the rate of food spoilage
- Increase the rate of product formation (pharmaceutical, industrial processes)
Units used to measure a rate of a reaction vary, for example:
Δ[]/Δt = mol/L*s (mostly) or g/s or mol/s or l/s
Rates are measured: Using the mole concept or Graphically by calculating the slope of a curve.
Method 1: Using the Mole concept to find the rate of a reaction: can be calculated in two ways: A –> B
- how fast a reactant disappears –> rateA = -Δ[A]/Δt (usually neg which is why neg is infornt to make it pos since rates are always expressed in a pos change of concentration over time)
- how fast a product appears –> rateB = Δ[B]/Δt
(can use mol ratios to figure out the rates for products and reactants in the same reaction)
Method 2: Graphically:
- Calculating average rates of reaction –> if decreasing on graph then reactant but if increasing then product –> average rate of reaction can be calculated from the slope of the secant on a concentration-time graph –> ratea/b = Δ[]/Δt
- calculating instantaneous rates of a reaction –> instantaneous reaction rate is the rate of a chemical reaction at a single point in time –> it can be calculated from the slope of the tangent on a concentration-time graph
Note: rates of reaction are always expressed as a positive value, the negative placed in front of the reactant species is to ensure that the two negatives multiply to give a positive value even though the reactant value calculated will be negative.
Stoichiometric rate relationships: in general aA + bB –> cC + dD
so rate = -1/aΔ[A]/Δt = -1/bΔ[B]/Δt= 1/cΔ[C]/Δt= 1/dΔ[D]/Δt
- stoich of the rxn will tell you that the rate of a appearance of something (product) is equal to half or whatever the number is of the rate of disappearance of something else (reactant)
- this is rate equation of rxn
Based on experimental data, it was found that it is not necessary for all reactants to have an impact on the rate of a reaction. Each type of reaction has its own unique path and which reactant(s) will impact the rate of the reaction. There is a strong relationship between concentrations of reactants and rate of reaction. All rates are based on the initial concentration.
Reaction rate depends on the reactant concentration raised to a power.
The relationship between reaction rates and reactant concentrations can be expressed in an equilibrium:
A + B --> Products
The rate of the reaction is:
rate is proportional to [A]^m[B]^n
replacing proportionality with a constant gives us: Rate law –> rate = k[A]^m[B]^n
- k = rate constant; value IS temperature DEPENDENT
- n,m - order of reaction; value can ONLY be calculated EXPERIMENTALLY
- Exponents (n,m) - order of reaction can have: Whole numbers (both negative and positive), Fractions and Zero values: which means the rate is independent of the concentration of that specific reactant.
(we already know that collision, temp, surface area, catalyst and agitation all affect the rates)
First Order Reaction: reaction rate depends on the concentration of only one species irrespective of the number of reactants present.
Rate = k [reactant]1
Second Order Reaction - (exponent =2)
The rate depends on the concentration of one species raised to the power of two OR it depends on two different reactants raised to the power one.
A 🡪 products A + B 🡪 products
Rate = k[A]2 Rate = k [A]1[B]1 (overall order is 2)
Overall reaction order - is the sum of exponents of the concentrations given in the rate law.
Rate = k [A]x[B]y Overall order of the reaction is (x+y) - rate is mol/l*s --> k units vary while [] is mol/l
Zero Order Reaction:
Rate = k[A]0 = k*1 = k
Rate = k ; therefore rate is constant, independent of reactant concentration.
Measuring reaction rates
Example:
MgCO3 + HCl 🡺
Collection of gas : monitor changes in volume of gas or changes in pressure in a fixed volume. Note: must ensure that the gas being produced in the reaction is not significantly soluble in the solvent being used.
Note: Must correct for pressure by taking vapour pressure into account.
Measure the change in mass of magnesium carbonate.
Measure change in pH via the consumption of HCl over time.
Measure changes in colour intensity using spectrophotometer or just visually.
Electrical conductivity: production of ions can be measured directly to current. The higher the concentration of ions present in solution the greater the electrical conductivity.
Clock reactions: in some reactions a product can be consumed by further reaction with another added substance. This added substance will react with the product which will result in a observable change (colour change). The time that it takes for this colour to be observed can be used to measure the rate of the reaction.
Titration: One can remove small samples from the reaction mixture at different times and then titrating the sample to determine the concentration of either one of the reactants or one of the products at this time. Use the data to graph concentration versus time. Usually works great with reactions that proceed slowly.
What about reactions that proceed really fast? Titration can still be used, but the reaction must be stopped by QUENCHING the reaction. So conditions of the reaction must be changed so that no more products are being formed at that time. For example, the temperature can be decreased rapidly or by adding an excess of a compound that will rapidly react with one of the reactants.
Example: CH3Br + OH- 🡪 CH3OH + Br-
In order to quench this reaction, excess amount of a strong acid can be added to consume the hydroxide ions.
Order of RXNS and rates
The rate of a reaction is defined as the change in the concentration of a reactant (or product) in a certain time interval.
rate equations are: rate = -1/aΔ[A]/Δt = -1/bΔ[B]/Δt= 1/cΔ[C]/Δt= 1/dΔ[D]/Δt
The combined affects by the reactants can be shown by: r = k [A]x[B]Y[C]Z
‘k’ is the rate constant: it is a constant characteristic of the reaction. It is the rate when [ ] of all reactants is 1 mol/L.
‘x, y, z’ are the order of the reaction with respect to the individual reactant (how it affects rate).
‘[A][B]…’ depends on how many reactants there are.
The rate law expression can only be determined through experimental data.
In a first order reaction k the reaction concentration changes in the same way as the rate.
r = k[N2O5]^1
if you double n2o5 the rate also doubles
In a second order reaction, the reactant may double while the rate quadruples.
r = k[CO]^0[NO2]^2
doubling CO causes no change in rate while doubling no2 would increase the rate 4 folds
So how do we experimentally determine the order of a reactant and the rate constant? Reactions are carried out at a specific temperature and the rate measured.
you can eyeball the data to figure out the order or you can use the general equation: rate 2/ rate 1 =( []2/[]1)^x –> plug in the given values, then solve and find the common base to get the value of the exponent x and the rate would be –> rate = k[]^x
if 2 different types of reactants first focus on one reactants order then the others. See where one is kept constant while the other is being change and see affect on rate. use the 2 experiment and do rate 1/rate = ([a]1/[a]2)^x * ([b]1/[b]2)^y –> solve for one of the variables to get order (usually one is kept constant meaning it would be 1^x which is always one). Then solve for the other experiment to find the other order.
To determine the rate constant, take one data set and substitute into the rate law equation to solve for k. Note: the units for the rate constant will change depending on the order(s) of reactants involved. –> do k = r/[a]^x etc to find k –> put in one experiments values to solve and then do unit analysis.
if multiple reactants try to see in which experiments is only one reactant varying and rest are constant to get the orders. if zero you can leave out when finding the rest.
if given a picture, visually identify or make a chart to find the orders (see where ones are constant and one is varying to get orders. will be given the rate)
FACTORS AFFECTING RATES OF REACTION
A chemical reaction is evident when the products have distinctly different properties from the reactants. Hence, a new substance has been made. In order for a new substance to be made, the forces of attraction in the reactants must be broken (requires energy). The atoms are rearranged and new bonds are formed from the separated particles (releases energy). However, a chemical reaction in not just dependent on having enough energy to break the initial attraction between the atoms of reactants. Collision Theory also plays an important role. In order for a chemical reaction to occur, the reacting molecules must collide with each other with a certain orientation. The rates of the reaction depend on the number of successful collisions between the reacting molecules.
FACTORS AFFECTING RATES OF REACTION:
Some chemical reactions take place almost instantaneously, whereas other take a long time. Being able to control the rate of reaction is of particular concern in industries. Knowledge in chemical kinetics is essential in being able to increase or slow down a reaction.
Temperature: increasing the temperature of a chemical reaction will almost always increase the speed of the reaction.
- Explanation: “ At higher molecular speeds there are more molecular collisions and hence a greater chance of a reaction. If the average molecular kinetic energy increases, more of the collisions will have enough energy to overcome the activation energy barrier.”
Surface area (size of solid particle): As size decreases the rate increases.
- Explanation: “With greater surface area, the collisions of other molecules with the surface are more frequent. The more collisions result in a faster reaction.”
Stirring/Agitation: As stirring increases the rate of the reaction generally increases as well.
- Explanation: “With increases in stirring, this ensures that reactants will come in contact with each other (increase in surface area).”
Phase of reactants: In general, reactions in solution, reactions of liquids and reactions between gases occur much faster than solid phase reactions.
- Explanation: “When in a solution, the ions are already separated. Therefore, the ions can react almost immediately.”
Concentration of reactants: Generally, an increase in concentration of a reaction will increase the rate. The greater the number of molecules present the higher the chance of collision occurring.
- Explanation: “the more concentrated the reactants, the closer they are to each other. Therefore, there is a greater probability of collision. Most collisions will result in a faster reaction.”
Catalysts: A substance that is used to speed the rate of the reaction but is not consumed during the reaction. A catalyst can be recycled due to the fact that it is not used up.
- Explanation: “It provides an alternate pathway with a lower activation energy. Therefore, at the same temperature, more molecules have enough energy to overcome a new, smaller activation energy barrier. Reaction will be faster.”
Collision Theory and Rates of RXN
Through experimental observations a rate can be determined. What happens to the rate when various changes are to it, for example: concentration, temperature, surface area and catalysts? These factors do have the ability to alter the rate of a reaction.
Recall: Molecules held together by chemical bonds.
The collision theory explains the reaction rates with respect to the fact that reacting particles (atoms, molecules, or ions) must collide with one another. In order for a reaction to occur, there is a need for an effective collision. An effective collision is one in which all the conditions are met such that products are formed from the
reactants. There are two criteria that need to be fulfilled in order for an effective collision to occur, these are:
1. Correct orientation of reactant (collision geometry)
2. Sufficient collision energy
Therefore, in order for a reaction to occur, the reacting molecules must be facing the direction for appropriate impact and the molecules colliding must have a certain energy to cause the bonds to break. This minimum energy required by reactant particles was described by Svante Arrhenius in 1889 as ACTIVATION ENERGY (E a ).
Rate frequency of collision x fractions of collisions (% E a barrier) x correct orientation
(to increase frequency of collions, increase [], s.a, t, and agitation. to increase fractions of collisions increase catalyst and maybe temp and for correct orientation you need a catalyst)
The magnitude of activation energy requirement depends upon the nature of the reactant (the kinds of chemical bonds that must be broken in the reactant if the reaction is to proceed). Therefore, each particular reaction will have its own particular set of activation energy for the reactants involved. Changes in
concentration, temperature and surface area will not have an affect on the activation energy, this will remain the same. However, the frequency of collisions will be affected (alters number of particles which will have enough energy to
collide effectively for a reaction).
Maxwell-Boltzman Distribution: graph that shows the number of particles on y axis and kinetic energy increase on x axis. The hump is where there’s an average. graph towards right means there is a temp increase and more k.e. if going towards left then a decrease in temp and low k.e.
Some important things to remember about collision theory:
1. Particles are always moving in random motion at various speeds. (like seen in the boltzman curve)
2. A chemical reaction must involve collisions between particles or with the walls of a container.
3. An effective collision is essential to make products. (right energ in the right orientation)
4. Ineffective collisions return to their original state (reactants) due to the lack of correct orientation or lack of energy to overcome activation energy
barrier.
5. The rate of a given reaction depends on the frequency of collisions and the fraction of those collisions that are effective.
Energy Profiles:
graph that shows the potential energy on y axis and rxn profile, progress and coordinate on the x axis. shows a curve from the reactants to products. From reactants to the height of the curve is Ea or activation energy (forward). At the height of the curve is []=/ which is a transition state. A.C is an activated complex –> it is unstable because of the octet breaking, had bonds breaking and reforming, is neither a reactant or product. enthalpy is also shown and is written as pos or neg.
At the transition state (top of the activation energy barriers) an activated complex exists. This activated complex is unstable molecule with a particular geometry, this is because the species possess maximum potential energy at this
stage. Therefore, there can be two fates following the activated complex stage. One is that the activated complex may return to reactants or continue to form product molecules.
Transition state theory: used to explain what happens when molecules collide in a reaction specifically at the transition state where the activated complex is formed.
there are potential energy diagrams or energy profiles for both exo and endothermic reactions. For Ear(reverse) of exo it is = |enthalpy|+ Eaf
for endo its Ear = Eaf - |enthalpy|
Activation energy and temp:
in the MB distribution curve the activation energy can be shown. molecules passed this have sufficient energy to overcome Ea. when temp is increased the graph is moved to the right and a greater number of particles can now overcome the Ea as they have enough energy. Experimental data shows that when temp increase the rate increase by 10x. 2 major factors/functions:
1. more collision frequency
2. flattening of MB curve allows for more molecules to overcome the Ea value.
therefore temp in the rate proportional equation affects the frequency of coliion and fractions of collision (Ea% barrier). this results in a more pronounced rate of rxn increase.
Activation energy and catalyst:
Catalysts: a substance that speeds up a reaction, but is chemically unchanged at the end of the reaction. When the reaction is finished, you would have exactly the same mass of catalyst as you at the beginning.
Examples of catalysts:
MnO 2 , concentrated sulfuric acid, Iron, Nickel, and Platinum.
Catalysts provide an alternative pathway for the reaction to happen with lower activation energy. By doing this, there will be an increase in the number of successful collisions.
in the MB curve with a ctatalust the Ea is lowered or moved towards the left so there are more particles that can actually overcome the Ea.
therefore catalyst affect the fractions of collions (%Ea barrier) and the correct orientation.
Catalysts: provide an alternative pathway with a lower activation energy (E a ) barrier. This allows molecules to react as compared to a reaction without a catalyst in a shorter amount of time.
So in the rate is proportional to frequency of collision * fraction of collisions (%Ea barrier) * correct orientation
- [] - frequency
- T - frequency * %Ea barrier
- S.A - frequency
- catalyst - %Ea barrier and correct orientation
- agitation - frequency
in the rate law equation k is affected by temp, catalyst, as and agitation
K = Ae ^ -Ea/RT
where Ae is y-nt and a frequency factor. RT would be slope?
take ln of both sides
lnK = -Ea/RT + lnA
(lnK is y, Ea/RT is m and lnA is b)
can be graphed (x axis is 1/t)
Rate Mechanisms
The rate law expression presents some interesting questions. For example, why does one reactant increase the rate of a reaction while
another may have no effect on the rate while, it is necessary for the
reaction to occur? Hint: think about the highway example.
Most reactions are made up of a series of steps called ELEMENTARY STEPS. Elementary steps make up the REACTION MECHANISM not a
single step. Depending on the reacting species, the chemical reaction can be a single step or many steps for a single reaction there can be many proposed mechanisms. Note: reaction mechanisms must be determined experimentally. For this reason, there are many opportunities for research in this area of chemistry.
Reaction intermediates: species that are formed during the course of the
reaction but immediately react again and are not present when the reaction is complete.
Molecularity: The number of reacting species:
- Unimolecular: collision of single molecule with container walls.
- Bimolecular: collision between 2 molecules.
- Termolecular: collision between 3 molecules. (Unlikely situation)
- Anything higher than 4 molecules colliding is a HIGHLY UNLIKELY
situation. Very rare in nature.
However, not all steps occur at the same rate. In actual fact, the slowest or RATE DETERMINING STEP determines the overall rate of a chemical reaction. Therefore, when an increase in concentration of a reactant does not affect the reaction rate, we conclude that the reactant is not involved in the rate-determining step of the reaction mechanism for the reaction.
What must be done to determine reaction mechanisms?
1. A rate law equation must be determined based on experimental data.
2. The species present in the rate law are the ones that react the slowest and thus determine the rate-limiting step.
3. Based on the rate law equation, a ‘probable’ mechanism can be made to accommodate the rate law.
4. Important: All elementary steps in the proposed mechanisms must always add to give you the actual reactant and product species that one starts out with.
Example: 4HBr (g) + O 2(g) 2H 2 O (g) + 2Br 2(g)
Note: This reaction involves the collision of 5 molecules to obtain the
products. This is a very unlikely situation. Thus there must be elementary step that allow for this reaction to occur.
The rate law equation (determined experimentally): r = k [HBr][O 2 ]
What is the order of the reaction? 2nd order
Based on the equation above it can be concluded that: slowest step is 1hbr + 1 o2 –> intermediate
*coefficent in the slow step come from the exponents in the rate law equation
mechanisms for reactions are proposed, intermediates can be made up. write if slow or fast. add and cancel for target equation.
all mechanisms can be shown on the energy profile showing that the first rate step or rate determining step has a huge Ea
Generally, the rate equation that is determined experimentally is
r= k [molecule X] m [molecule Y] n
Then the rate-determining step in the mechanism must be:
mX + nY products or reaction intermediate
