This document describes an experiment to measure the rate of reaction between magnesium and sulfuric acid. The rate is measured by collecting hydrogen gas produced over time. The results show the volume of hydrogen increases rapidly at first as the reaction proceeds, then slows down as the magnesium is used up. The document also discusses factors that generally affect chemical reaction rates, such as concentration, surface area, temperature, and catalysts.

1. Tuesday 14 February 2017

2.  Chemical reactions take place at very different
rates – some are very fast (like explosions) and
others may take months or years to proceed.

3.  In this experiment the reagents are magnesium
and sulphuric acid.
 The products are magnesium sulphate and
hydrogen gas.
 The rate of the reaction is measured by measuring
how fast the hydrogen gas is produced.
 The limiting reagent is the amount of magnesium.
This is the first reagent to be use up.

5. Rate of reaction experiment
Time (s) Volume of Hydrogen (ml)
0 0
30 16
60 25
90 31
120 34
150 35
180 36
210 36
240 36

6.  We can record the rate using this relation|:
 Rate of reaction = change in recorded property
time for the change
 Actual rate is the gradient of the line of the product
concentration versus time graph.
 Average rate is the gradient of the line joining the two
point in time over which the rate is being measured.

7. Rapid
reaction here
Reaction slows
down here.
Reaction
stops here

8. Rate at the start of
the reaction.

9. Average rate of
reaction between 0s
and 200s.

10.  In this experiment calcium carbonate reacts with
hydrochloric acid.
 The rate of reaction is measured by collecting the
carbon dioxide gas produced in a gas syringe.

11.  Calculate the number of moles of hydrogen
produce in experiment 1 above.
 Volume of hydrogen = 36ml
 1 mole of hydrogen = 24000 ml
 Number of moles of hydrogen = 36 ÷ 24000
 = 0.0015 moles

13.  In this experiment calcium carbonate reacts with
hydrochloric acid.
 The rate of reaction is measured the mass of the
remaining chemicals. The loss in mass is due to
the escaping carbon dioxide gas.

15.  In this experiment sodium thiosulphate reacts with
hydrochloric acid to produce sulphur.
 The rate of reaction is measured by timing how
long it takes for a cross drawn below the beaker to
disappear. The cross should disappear once the
same amount of sulphur is produce in each trail.

17.  Concentration
 The higher the level of concentration (or pressure
in gases) the faster the reaction.
 This is due to increased collisions between
reacting particles.

18.  Surface area in solids
 The larger the surface area of a solid the faster
the reaction. Finely divided substances have
much larger surface areas than large chunks of a
solid.
 This is due to increased collisions between
reacting particles.

19.  Temperature
 The higher the temperature the faster the reaction.
 This is due to increased and more energetic
collisions between reacting particles.
 A 10o
C rise in temperature often results in a
doubling of the reaction rate.

22.  Only molecules with enough activation energy will
react to form the products.

23.  Catalysts
 Catalysts increase the rates of chemical reactions.
 This is due a lowering of the activation energy for
the reaction.
 Catalysts are not consumed in the reaction so in
theory they can be used over and over again. In
practice they are often contaminated and/or some
is lost in a process.

26.  Scientists assume all gases are made of particles in
constant random motion. They have regular elastic
collisions with other molecules.
 Gas molecules collide with the walls of their container
and exert pressure but do not lose energy in their
collisions and do not attract other molecules.
 The volume of actual gas molecules in a container is
negligible and their average kinetic energy is
proportional to the temperate (Kelvin).