Week 9 - Kinetics (activation energy, elementary reaction steps, reaction rate, catalysis) Flashcards
What is a catalyst? (5)
A catalyst is a substance that increases the reaction rate without itself being consumed in the reaction.
In general, a catalyst provides an alternative reaction pathway that has a lower total activation energy than the uncatalysed reaction.
A catalyst will speed up both the forward and the reverse reactions.
Added in non-stoichiometric amounts (~ 10-6
to 10-1X)
A catalyst does not affect either △H or the overall yield for a reaction.
Catalysis – why is it important?
- More than 90% of industrial processes actually use catalysts in one form or the other.
- If you are employed as a materials scientist, it’s likely that you will encounter tasks including catalysis.
Who uses catalysts? (4)
- Petroleum and energy production
- Chemicals and polymer production
- Pollution control
- Pharmaceutical and food industry
Activation Energy
energy required to overcome the reaction
barrier. Usually given a symbol Ea or ∆G^≠
- The activation energy (Ea) determines how fast a reaction occurs, the higher the activation barrier, the slower the reaction rate. The lower the activation
barrier, the faster the reaction.
Activation energy: catalyst effect
Catalyst lowers…
the activation energy for both forward and reverse reactions, and consequently the catalyst increases the rate of reaction.
Importance of catalysis: (6)
- increasingly important in synthesis
- selectivity in production of fine chemicals
- clean processes, high atom economy (bulk processes)
- production of high-tech products / materials
- mild conditions (low energy consumption)
- environmentally friendly
Catalytic reaction simple process steps:
Every catalytic reaction is a sequence of elementary steps, in which reactant molecules bind to the catalyst, where they react, after which the product detaches from the catalyst, liberating the latter for the next cycle.
Catalytic activity
is the increase in the rate of a specified chemical
reaction caused by a catalyst under specified conditions
Catalyst selectivity
the ability of the catalyst to drive the reaction
towards a particular product, as there may be more than one in typical catalytic reactions
Catalyst stability
refers to the chemical, thermal, and mechanical
stability of a catalyst, which determines its lifetime in industrial reactors.
The types of catalysts,
Classification based on the its physical state, a catalyst can be:
- gas
- liquid
- solid
Different types of catalysts,
Classification based on the substances from which a catalyst is made:
- Inorganic (gases, metals, metal oxides, inorganic acids, bases etc.)
- Organic (organic acids, enzymes etc.)
Different types of catalysts,
Classification based on the ways catalysts work:
- Homogeneous - both catalyst and all reactants/products are in the same phase (gas or liquid)
- Heterogeneous - reaction system involves multi phase (catalysts + reactants/products)
Different types of catalysts,
Classification based on the catalysts’ action:
- Acid-base catalysts
- Enzymatic
- Photocatalysis
- Electrocatalysis, etc
Bio-catalysis
- Biocatalysis is the chemical process through which enzymes or other biological catalysts perform reactions between organic components.
- Reactants and catalyst may be in the same phase or in different phases.
- Michaelis-Menten
Biocatalysis: Pros & Cons
- Very high selectivity; stereo-selectivity, regio-selectivity (allows selective modification of a specific site in a molecule), and functional group selectivity.
- Operational advantages; reactions under mild operational conditions.
- Catalysts that are biodegradable thus typically “greener” and more sustainable.
- Often very costly
- Requires a development timeline that is too long to meet the needs of real world industrial manufacturing
Homogeneous Catalysis
A homogeneous catalyst exists in the same phase as the reactants.
- An example is the decomposition of hydrogen peroxide in aqueous solution
The Catalyzed Decomposition of H2O2:
- A small amount of NaBr is added to a solution of
H2O2 (A). - Oxygen gas forms quickly as Br− (aq) catalyzes the H2O2 decomposition; the intermediate Br2 turns the solution orange (B).
- The solution becomes colourless when the Br2 is consumed in the final step of the mechanism (C).
Homogeneous catalysts: pros and cons
- Catalytic mechanisms are relatively easy to study in homogeneous systems, using powerful methods as NMR to both assign structures
and follow reaction kinetics. - High selectivity.
- Homogeneous catalysts have the disadvantage that they can be difficult to separate from the product.
- Poor stability
Heterogeneous catalysis
Heterogeneous catalyst are catalysts for which the catalyst and the substrates for the reaction are in different phases, cf. solid catalyst
and gas-phase reactions.
- A heterogeneous catalyst is in a different phase than the reaction mixture, e.g hydrogenation, the addition of H2 to C═C bonds to form C―C bonds
Simplified process: absorption, reaction, desorption
Heterogeneous catalysis: pros and cons
- High catalyst activities
- Cheap
- Ease of preparation
- Often low selectivity
- Deactivation
Application of catalysis,
Industrial applications:
Almost all chemical industries have one or more steps employing catalysts:
Petroleum, energy sector, fertiliser, pharmaceutical, fine chemicals …
Advantages of catalytic processes
- Achieving better process economics and productivity
- Increase reaction rates - fast
- Simplify the reaction steps - low investment cost
- Carry out reaction under mild conditions (e.g. low T, p) - low energy consumption
- Reducing wastes
- Improving selectivity toward desired products - less raw materials required, less unwanted wastes
- Replacing harmful/toxic materials with readily available ones
- Producing certain products that may not be possible without catalysts
- Having better control of process (safety, flexible etc.)
- Encouraging application and advancement of new technologies and materials
And many more …
What is Chemical Kinetics:
study of reaction rates,
and the factors that affect them.
What is Reaction Rate:
the change in the concentrations of reactants or products as a function of time
Rates of reaction vary greatly for everyday processes
Kinetics vs thermodynamics
(what’s the difference)
Thermodynamics
* Looks at the starting and end points of a
reaction.
* Predicts the direction of spontaneous
reactions.
* Predicts what the final concentrations will be
if equilibrium is reached.
* The variable of time is irrelevant in
thermodynamics.
* The rate or mechanism of chemical reactions
are not considered
Kinetics
* Cannot predict the direction or end point of a
spontaneous reaction.
* Looks at the rate of chemical reactions.
* Analyses the reaction mechanism, i.e. the
pathway the reaction takes.
Why do reactions that are thermodynamically
possible sometimes proceed slowly or not at all?
In chemical reactions a barrier needs to be overcome as chemical bonds need to be broken and new ones need to be formed.
What determines whether a reaction is exo or endothermic?
(based on delta H)
Thermodynamics (for ΔG spontaneity, for ΔH,
exothermic or endothermic) only depends on
how reactants’ and products’ energy compares,
not on how the transformation takes place!
Activation Energy
- In order for the reaction to occur, bonds in the reactants need to be broken so new bonds, and thus the products, may form.
- The energy that is needed for the ‘old bonds’ to be broken is the activation energy (Ea).
Reaction mechanism:
What does activation determine?
Ea determines how fast the reaction proceeds: think of it like an energy “mountain” that the
reactants need to climb before they can react.
EACH “pathway” will have its own Ea. The one with the lowest activation energy is the fastest so it “wins” (but if the activations energies are similar, you may end up with a mixture of products, each coming from a different pathway).
In chemistry, these pathways are called: mechanisms (reaction mechanisms).
Reaction mechanism
As you provide energy to the reactants those that collide, come close together and if the energy provided is the same or larger than Ea
(and the collision is in the right orientation), these
reactants colliding can form a transition state:
a short-lived chemical species where “old bonds” are partially broken and “new bonds” are partially formed. Short lived means VERY small fractions of a second, and they cannot be isolated.
Transition state:
The peak of the reaction pathway in an energy diagram, Gibbs Free Energy on the y and Reaction Coordinate on the x
delta G ^ (equals sign with perpendicular line going through it) - Change in Gibbs Free Energy
Elementary reactions:
- A reaction where a single energy barrier has to be overcome.
- Many reactions proceed through a number of elementary reactions.
In some cases there are two activation energies, and two transition states (one per step), and an intermediate: Don’t confuse intermediate with transition state!
Reaction Intermediates:
What’s the difference between Intermediates and transition states?
Intermediates:
* An intermediate may be “short-lived” (very
reactive) or not, but POTENTIALLY it could be
isolated, and it can be “detected” or characterised
by spectroscopy.
* Intermediates have full bonds (or fully broken
bonds, like incomplete octet species such as
cations or free radicals).
* Intermediates represent a stop in between
reaction steps.
* Reaction mechanisms are characterised by the
type and number of the elementary reactions
involved and the appearance of intermediates
Transition state
* transition states are VERY difficult to characterise (let alone isolate…)
* transition states have partial bonds being formed or broken.
* Transition states are NOT a “stop” (the reaction
proceeds through the transition state but NEVER
stops there)
Collision theory:
Collision theory states that two (or more) reactant species need to collide with the right energy and geometry for a reaction to take place
This means that if we increase the probability of collisions, the probability of “right” collisions, and therefore the rate of the reaction, will increase
Increasing the concentration of reactants will increase the rate
Increasing the energy of the system will increase the rate
Rate constant & activation energy
- Arrhenius equation
Rate constant:
𝑘 = 𝐴e ^ (-Ea / RT)
Rate equation, we will see more in detail but 𝑟 ∝ k
In general, rates of reactions increase as concentrations increase since there are more collisions occurring between reactants.
The overall concentration dependence of reaction rate is given in a rate law.
For our general reaction: l𝜈𝐴l 𝐴 + l𝜈𝐵l 𝐵 → l𝜈𝐶l 𝐶 + l𝜈𝐷l 𝐷
The rate law might look like:
r= k . [A]|^m . [B]^n
-The exponents m and n are called “reaction orders”.
-The reaction is mth order in A and n th order in B.
-The proportionality constant k is called the rate constant.
-The overall reaction order is the sum of the reaction orders:
m + n
Note that the reaction order is not necessarily related to the stoichiometric coefficients of the
reactants
How are Rate laws determined?
Rate laws are always determined experimentally.
How is reaction order always defined:
Reaction order is always defined in terms of reactant (not product) concentrations.
The order of a reactant is not related to the stoichiometric coefficient of the reactant in the balanced chemical equation.
Rate Law
Working out the overall reaction order
The rate law might look like:
r= k [A]^m [B]^n
-The exponents m and n are called “reaction orders”.
-The reaction is mth order in A and n
th order in B.
-The proportionality constant k is called the rate constant.
-The overall reaction order is the sum of the reaction orders:
m + n
Note that the reaction order is not necessarily related to the stoichiometric coefficients of the
reactants.
In rate laws what does [ ] mean?
concentration
What is the rate law
In materials chemistry, the rate law typically refers to the mathematical expression that describes the rate of a chemical reaction in terms of the concentrations of reactants. The rate law helps in understanding how the rate of a reaction depends on the concentrations of the reactants involved.
In rate laws when plotting a graph, its always…
conc on the y
and time on the x
so then delta conc is being measured
Explain what is meant first, second and zeroth order
If we keep everything the same but double the conc of one reactant, and if as a result the rate doubles then the reaction is first order, with respect to that reactant
instead if we double the conc and the rate quadruples, then the reaction is second order, and if a change doesn’t effect the rate then its zeroth order.
The overall reaction order just the sum of the individual orders.
When we double a conc, what happens to the rate?
If the rate is multiplied by:
1
2
4
1/2
then the reaction order in respect to that reactant are:
0 (zeroth order)
1 (first)
2 (second)
-1 (order of negative one)
Rate Laws
If we want to discern the conc at anytime, what do we use?
Integration rate laws, we will use different integrated rate laws dependant on the reaction order.
Refer to image taken, on the laws that are in the format y = mx + c
Where x is always time (t) and m is either k or -k
