Chemistry - Chapter 16: Rates of Reaction

rate of a chemical reaction

the change in concentration per unit time of any one reactant or product

average rate: total volume of oxygen given off (cm³) / time for reaction to complete (secs) x 100

or unit of conc/unit of time

R = 1/t

instantaneous rate of reaction

the rate of reaction at any one particular time during the reaction

Collision Theory

for a chemical reaction to occur, the reacting particles must collide with each other with enough energy to produce a product. Energy required is called the activation energy (insufficiency results in no reaction)

activation energy

the minimum amount of energy that colliding particles must have in order for a reaction to occur

effective collision

one that results in the formation of products

large activation vs small activation energy required

large AE required → small no. of molecules have enough energy to pass over the barrier per second (slower reaction rate)

small AE required → faster rate

factors: nature of reactants

• ionic reactions = fast + covalent reactions = slow

• ionic → don’t involve bond breaking but simply the union of oppositely charged species at room temperature (precipitate forms quickly)

• covalent → reactions that involve the breaking of covalent bonds at room temperature (precipitate forms slowly)

• é.g. test for Cl-

factors: surface area/particle size

• the larger the particle size, the slower the rate of reaction due to smaller surface area

• more exposed particles → more frequent and successful collisions

• the smaller the particle size, the faster the reaction

• e.g. powdered CaCo3 vs marble chips

factors: concentration

• the no. of moles of a given substance in a certain volume of solvent

• an increase in concentration results in more particles per unit volume, hence more frequent collisions between particles

• changing the conc. of reactants in a chemical reaction alters the rate of reaction (HCl vs sodium thiosulfate → black cross on paper experiment)

• remember that there’s a maximum number of possible collisions between molecules and once you hit that maximum, you can’t get faster, no matter how much you put in

• concentration and rate are proportional to each other

factors: temperature

• temperature is proportional to the kinetic energy of particles. The more energy the particles have, the faster they move around

• increase in temperature → higher kinetic energy, leading to more frequent collisions & successful ones as collisions themselves occur with more energy

• rate and temperature are not directly proportional

• @ rough, approx. 10 degrees C increase doubles the rate of reaction

factors: catalysts

alters the rate of a chemical reaction without being consumed in the process of the reaction (provides alternative route w/ lower activation energy)

properties of a catalyst

• recovered chemically unchanged at the end of an experiment but may be changed physically

• tends to be specific (e.g. enzymes)

• only needs to be present in small amounts

• catalytic poisons → action of catalyst may be destroyed by these. e.g. platinum, palladium, rhodium

• in equilibrium reactions, catalysts help the reaction achieve equilibrium more quickly…doesn’t affect equilibrium’s position or final composition of mixture in the reaction → they are reversible in their action

homogenous catalysis

both the reactants and catalyst are in the same phase

é.g. 2H2O2 → 2H2O + O2

heterogenous catalysis

reactants and catalysts are in different phases

é.g.

• Haber process for synthesis of ammonia; reactants are gases and catalyst is iron, a solid

• oxidation of methanol to methanal over heated platinum. this is the experiment with the cyclic spiral glowing and cooling due to exothermic nature of the reaction (oscillation reaction)

autocatalysis

the catalyst is the product of the reaction

(e.g. Mn7+ + Fe3+ → Mn2+ + Fe2+, where Mn2+ is produced in the reaction and the rate increases dramatically)

immediate colour change of purple to colourless is observed

mechanisms of catalysis: surface adsorption theory

• explains most heterogenous catalysis

• adsorption - the accumulation of substances at the formation of another substance

• e.g. ethene & hydrogen

If a solid catalyst is placed among gaseous reactants, the gases adsorb onto the surface of the catalyst.

• There is now an increased concentration of the gases on the surface of the catalyst so more effective collisions can occur to form the product of the reaction.

• The product then leaves the surface of the catalyst and this leaves more room for more reactants to adsorb onto the surface again.

As mentioned before, the more finely divided the catalyst, the greater the surface area for adsorption to occur and so the faster the rate of reaction.

If a catalyst poison is present it takes up space on the surface of the catalyst and so the catalyst is useless.

An example of heterogeneous catalysis is the Haber Process. The reactants are hydrogen and nitrogen and the catalyst is finely divided iron. Another example is that of platinum catalysing the reaction between hydrogen and oxygen to form water.

mechanisms of catalysis: intermediate formation theory

Catalyst and reactant combine to form an unstable intermediate. These intermediates decompose forming products and regenerating the catalyst.​

so…

• thís intermediate is formed very quickly and decomposes as soon as it is formed

• é.g. A + B → AB (slow reaction)

• Add catalyst:

A+C → AC (fast) and AC + B → AB + C (fast)

é.g. iodine snake

Add iodide ions with H2O2:

H2O2 + I- → H2O + IO- (fast, where IO is the intermediate)

then

H2O2 + IO- → H2O + O2 + I- (fast)

use MnO2

é.g. 2:

When potassium sodium tartrate is oxidised by hydrogen peroxide using cobalt (II) as a catalyst the following is observed.

• Potassium sodium tartrate is dissolved in hot water.

• Cobalt (II) ions are added. ions are pink in colour.

• Hydrogen peroxide is now added and a colour green is observed. The green colour is the colour of the intermediate compound.

• Carbon dioxide and steam are given off (very vigorous reaction!).

• After a while, when the reaction is finished, a pink colour appears again. This shows that the Co2+ ions have not been used up in the reaction.

phase

a specific state of matter that implies a boundary

heterocatalysis example

Oxidation of methanol by a platinum catalyst:

• The hot spiral cools in transfer to conical flask then glows hot again once inside → platinum glows due to exothermic reaction

• popping sound → hydrogen gas exploding

• smell of methanol → methanol is oxidised

• spiral ceases glowing then heats up again → CO2 is produced due to incomplete combustion, poisons the catalyst, then leaves

catalyst

a substance that alters the rate of a chemical reaction but is not consumed in the reaction

catalytic converters

• a device in vehicles that converts pollutants to less harmful substances in the exhaust pipe of cars

• cars produce NO, CO, NO2, all of which are dangerous to the environment

• catalytic converters contain a metal (Pt, Pd, Rh) which convert using the surface adsorption theory

catalytic poison

a substance that makes a catalyst inactive

exothermic reaction

products that have less energy than the reactants = heat is given out

endothermic reaction

products have more energy than the reactant = heat is taken in