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NOTES ON KINETICS

Kinetics is the branch of chemistry which studies reaction rates (how fast) and mechanisms (how).

A reaction rate can be expressed as:

Rate = ∆concentration/∆time

Consider the reaction: A + 2B à AB2

Note that the rate of disappearance of reactants and rate of formation of product are directly proportional to the stoichiometric coefficients in the balanced equation.

If reactant A is consumed at a rate of 25.0 mole/L, at what rate is product AB2 produced?

If Reactant A is consumed at a rate of 45 mol/L, at what rate is reactant B consumed?

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Let’s draw a graph of how the rate of the reaction changes over time. We will start with all reactants and no products initially. Graph:

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• The tangent at any point on the curved line = slope = rate; that is• ∆Y/∆X = ∆ concentration/∆time = rate• the rate (slope) at the beginning of the reaction is greater than the rate (slope) later on

Mechanism

How do chemical reactions actually occur? Chemists have proposed a theory (or model) that all net chemical reactions actually occur as a series of individual bimolecular collision reactions called elementary reactions. These elementary reactions make up the steps in a reaction mechanism and add up to a net reaction. For a bimolecular reaction to be effective (that is, one that results in a product), two conditions must be met:

1. the two reacting molecules must have sufficient kinetic energy to overcome electron cloud repulsions as they approach each other.2. the two reacting molecules must hit at the correct angle so the “active sites” on each molecule can interact

The above two conditions will occur with greatest probability when two molecules collide (i.e. a bimolecular collision). Statistically, three molecules meeting the above two conditions at the same time is extremely rare.

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Let’s consider the following reaction:

2A + B + C à Z

According to the Collision Model this reaction would statistically never occur – four molecules all having sufficient KE and all hitting at the correct angle at the same time is practically impossible. But we do observe the reaction occurs! As good chemists let’s propose a mechanism by which this reaction can occur. Keep in mind that each step in our mechanism must consist of a bimolecular collision (statistically more likely to occur) and that the steps must add up to the net reaction observed.

Rate product formed

Step 1: A + B à X 106/sec

Step 2: X + C à W 102/sec

Step 3: W + A à Z 106/sec

What is the net reaction? Is this consistent with the observed net reaction?

What is the rate­determining step?

What is the overall (or net) rate at which Z is produced? Why?

What is an intermediate?

What is a catalyst?

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Factors Affecting the Rate of a Reaction – key idea: frequency of collisions

1. Nature of Reactants – since bonds may be broken or formed in a chemical reaction, the nature of the bond is an important factor affecting reaction rate:

• reactions that involve negligible bond rearrangement are usually rapid at room temperature – e.g. the reaction between ionic substances in aqueous solution may take place in a fraction of a second:

AgNO3(aq) + NaCl(aq) à AgCl(s) + NaNO3(aq) very fastclear clear white ppt clear

AgNO3(s) + NaCl(s) à AgCl(s) + NaNO3(s) forget it

Since the dissolving process has destroyed the bonds in the ionic crystal, there are no bonds (except weak electrostatic IMFA) to be broken in this reaction. All that is needed is a rearrangement of ions, and this occurs rapidly.

• reactions that involve the breaking of bonds tend to be slow at room temperature:

e.g. 2H2O2(l) à 2H2O(l) + O2(g) hν

e.g. 2KClO3(s) à 2KCl(s) + 3O2(g) heat

2. Concentration – an increase in the concentration of one or more of the reactants usually but not always increases the rate of the reaction. An increase in the concentration of a reactant increases the FOC.

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1. Temperature – if temperature increases, the KE increases – molecules move faster and FOC increases resulting in an increase in rate. This can be represented graphically (Ea is the activation energy the minimum energy required to start a reaction):

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4. Surface Area – the greater the surface area the greater the FOC.

5. Mechanism – those mechanisms that are complex and have very slow rate­determining steps will have slow rates.

6. Pressure – only effective for gases; an increase in pressure on a gases will decrease the volume (Boyle’s law) – this in turn increases the concentration which increases the FOC and rate.

7. Catalysts – a catalyst increases the rate of a reaction by providing an alternate mechanism – a mechanism whose slow step has a lower activation energy. A catalyst itself participates in the reaction mechanism but is always regenerated in the end.

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demonstration: catalytic oxidation of tartaric acid

H2C4H6O6(aq) + H2O2(aq) + CoCl2(aq) à [Co­complex clear clear pink Intermediate]

green

+ CO2(g) + H2O(l) + CoCl2(aq) pink

demonstration: elephant’s toothpaste (catalytic peroxide decomposition)

step 1: H2O2(aq) + I­(aq) à H2O(l) + IO­(aq)

step 2: H2O2(aq) + IO­(aq) à H2O(l) + O2(g) + I­(aq)

net:

identify any intermediates in this two­step mechanism:

identify the catalyst in this mechanism:

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Potential Energy Diagrams

When two molecules approach each other in a collision, the electron clouds experience a gradual increase in their mutual repulsions, and the molecules begin to slow down. As this happens, the kinetic energy of the molecules is gradually converted to potential energy. If the pair of colliding molecules had little kinetic energy to begin with, they come to a stop before their electron clouds have penetrated each other very much, and then they simply fly apart, chemically unchanged. When two fast moving molecules collide, they have a lot of kinetic energy that can be converted to potential energy. This means they are able to overcome the forces of repulsion between their electron clouds and their electron clouds may interpenetrate. If the angle of collision is correct, a reshuffling of valence electrons may occur. Old bonds may be broken and new ones formed resulting in a net chemical change occurring.

There is a minimum kinetic energy that reactant molecules must jointly possess to overcome the repulsions between their electron clouds when they collide. This kinetic energy, which is changed to potential energy at the moment of impact, is called activation energy (Ea). This activation energy acts as a potential energy barrier between the reactants and products.

A potential energy diagram of a reaction shows the relationship between the activation energy and the energy absorbed or released during a reaction. The horizontal axis is labeled “reaction coordinates” and represents the progress of the reaction. The vertical axis is labeled “potential energy” and shows the relative potential energies of the reactants, products, and activated complexes (intermediates).

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Note the following:

• The point from reactant to activated complex (intermediate) is the activation energy• For any exothermic reaction the PE of the products is less than the PE of the reactants• For any endothermic reaction (reverse of exothermic) the PE of the products is greater than the PE of the reactants• A catalyst lowers the activation energy of both forward and reverse reactions• A catalyst has no net effect on the ∆H of the reaction• Exothermic reactions have lower activation energies compared too the endothermic direction

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