rates of reactions

rate of reaction

is the change in concentration in unit over time of any one reactant or product, rate = change in concentration/ time taken

average rate calculations

total change in reactant or product/ total time

instantaneous rate

is the rate at a particular point in time during the reaction

balanced equation of decomposition of hydrogen peroxide

2H2O2 → 2H2O+O2

powdered oxide

has a larger surface area allowing for more contact with reactant molecules

controlling rates of reaction

being able to control the speed of chemical reactions is important in both in everyday and when making new materials on an industrial scale.

collision theory

not all collisions result in a reaction, for a collision to be successful the molecules must have a minimum amount of energy called the activation energy (Ae)

effective collision

a collision that results in the formation of a product

example of fast reactions

fireworks, silver nitrate and sodium chloride

examples of slow reactions

nail rusting, newspaper turning yellow

factors affecting rate of reaction

temp, concentration, particle size, catalyst, pressure

reasons why these factors affect

each has its unique reason, more collisions, more effective collisions per unit time

temperature

temp increases particles move faster allowing for the number of collisions to increase

particle size

the greater the surface are the more collisions per unit time which equals more effective collisions per unit time

concentration

increased number of particles increases number of collisions

pressure

increased pressure decreases volume

catalysts

provides an alternative pathway with a lower activation energy

activation energy Ae

minimum energy required for a reaction to occur

heat of reaction (dalta H)

difference between the energy of the reactants and the energy of the products

endothermic reaction

heat is taken in from surroundings and the products formed have more energy than the reactants

exothermic reactions

heat is lost to the surroundings and the products formed have less energy than the reactants.

catalysts meaning

speeds up a reaction and isnt used up in the reaction, they create an alternative pathway, lowers Ae

properties of catalysts

recovered chemically unchanged, specific - only work on certain reactants, only needs small amounts

catalytic mechanisms

surface adsorption theory, intermediate formation theory

surface adsorption theory

heterogenous catalysis when reactants and catalysts are in the different phases

steps to surface adsorption theory

catalysts adsorb reactant particles onto its surface, bonds are weakened reducing the Ae required to form product, product is released and catalyst remains unchanged

steps - intermediate formation theory of catalysts

catalyst combines with reactant, forms an unstable intermediate compound, intermediate decomposes to give products, catalyst is regenerated

intermediate formation theory of catalysts

homogenous catalysis reactant and catalyst in same phase

experiment; the effect of a catalyst on the rate of reaction

hydrogen peroxide naturally decomposes into water and oxygen, this reaction is very slow therefore we add a catalyst manganese dioxide (black powder)

experiment; effect of concentration on reaction rate

1 100cm3 of sodium thiosulfate 2 10cm3 of HCL once this is added to conical flask start timer 3 swirl flask and place it on filter paper with a black x drawn on it 4 record time taken for x to disappear 5 repeat using 80, 60, 40, 20 cm3 of Na2S2O3

experiment; effect of temperature on rate

same volume as concentration, place on hotplate, start stopwatch, use thermometer, remove conical flask + place on filter paper with x, stop timer once x disappears