Download 9.10 The Speed of a Reaction notes

RATES OF REACTION
Increasing the rates of chemical reactions is important in industry because it helps
to reduce costs.
The rate of reaction is the speed at which a chemical change takes place. It is
followed by measuring the rate at which reactants are used up or the rate at which
products are formed. This allows a comparison to be made of the changing rate of
a chemical reaction under different conditions.
Rate = amount of reactant used
time
or
Rate = amount of product formed
time
Chemical reactions can occur only when reacting particles collide with each other
with sufficient energy. The minimum amount of energy which particles must have
to react is called the activation energy. Reactions are fastest at the start
because this is when there are the most particles. As a reaction proceeds,
reactants are used up and the rate decreases.
Concentration
of reactant
Time
Since the rate of a reaction is a measure of how quickly reactants are being used
up, it follows that we can determine the rate of a reaction by measuring the
gradient of the above graph.
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Various factors alter the rate of a reaction:
 the state of division of the reactants (i.e. the particle size of a solid): the
smaller the particle size, the faster the reaction is.
 the concentration of dissolved reactants or the pressure of gases: the
higher the concentration or the pressure, the faster the reaction is.
 the temperature of the reaction mixture: the higher the temperature, the
faster the reaction is.
 the addition of a catalyst speeds up the reaction
1. The Particle Size of Solid Reactants
The reaction used to study the effect of particle size is the reaction of calcium
carbonate, in the form of marble, with dilute hydrochloric acid.
CaCO3(s) + 2HCl(aq)
CaCl2(aq) + H2O(l) + CO2(g)
The reaction is followed by the change in mass of the reaction flask with time as
carbon dioxide is given off. It would also be possible to measure the change in the
volume of carbon dioxide given off with time by collecting the gas in, for example,
a gas syringe.
Method:
A constant mass of marble chips was weighed out and placed into a 250cm 3 conical
flask. 100cm3 of 1M hydrochloric acid was added from a measuring cylinder. The
flask was loosely stoppered with cotton wool (to allow the gas to escape but to
prevent the loss of liquid splashes) then placed onto an electronic balance. The
mass of the conical flask was recorded every 15s for the first minute, then every
30s for a total of ten minutes. It is assumed that the temperature of the reaction
mixture stayed constant.
The experiment was repeated with three different sizes of marble chips, keeping
all other variables the same. The results were tabulated and a graph of mass of
carbon dioxide (y-axis) against time (x-axis) was plotted.
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Results:
Large Particles
Time /s
Mass of
flask/g
Mass of
CO2 /g
Medium Particles
Mass of
flask /g
0
15
30
60
90
120
150
180
210
240
270
300
330
360
390
420
450
480
510
540
570
600
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Mass of
CO2 /g
Small Particles
Mass of
flask /g
Mass of
CO2 /g
Interpretation of Results:
The reaction occurs only on the surface of the calcium carbonate. For a given mass
of calcium carbonate, the smaller the size of the particles, the greater is the
surface area and the faster the reaction. This is shown on the graph by the
gradient becoming steeper as the particle size decreases.
The smaller the particle size (the bigger the surface area) , the
faster the reaction
Since the same quantities of reactants are involved in all three reactions, the same
mass of carbon dioxide is given off in all of them, if the reaction is allowed to go to
completion. This is shown by all three curves levelling off at the same total mass of
carbon dioxide.
The calcium carbonate is in excess, so the hydrochloric acid is used up completely
as the reaction takes place. Since its concentration decreases with time, the
reaction becomes slower and slower. This shown on the graph by a curve of steadily
decreasing gradient.
The expected mass of carbon dioxide can be calculated. Since the calcium
carbonate is in excess, the mass of carbon dioxide depends on the amount of
hydrochloric acid used. 100cm3 of 1M HCl contains 0.1mole of HCl.
CaCO3(s) + 2HCl(aq)
2 moles
0.1 mole
0.1 mole
0.1 mole
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CaCl2(aq) + H2O(l) + CO2(g)
1 mole
0.05 mole
0.05 x 44 g
2.2g
4
2. The Temperature
The reaction used to study the effect of temperature is the reaction of sodium
thiosulphate solution with dilute hydrochloric acid.
Na2S2O3(aq) + 2HCl(aq)
2NaCl(aq) + S(s) + H2O(l) + SO2(g)
The reaction is followed by the appearance of colloidal sulphur as the reaction
proceeds. The formation of sulphur begins as soon as the reactants are mixed, but
it takes time for observable amounts to be produced. The time taken to reach a
particular point in the reaction can be determined by standing the reaction flask on
a piece of paper marked with a feint cross and timing how long it takes for the
cross to be obscured when looked at from above.
Method:
10cm3 of sodium thiosulphate stock solution (concentration 40g per litre) were
measured out into a measuring cylinder and poured into a 250cm 3 conical flask.
40cm3 of water were measured out in a similar way and added to the flask. The
flask was heated gently over a Bunsen burner to a temperature slightly above the
desired temperature. The flask was placed on a piece of paper marked with a feint
cross. 5cm3 of 2M hydrochloric acid were measured out into a second measuring
cylinder. As the acid was poured into the conical flask, a stopwatch was started
and the mixture gently swirled. The initial temperature was recorded. The
stopwatch was stopped when the cross, viewed from above, became obscured. The
time and the final temperature were recorded. The mean of the initial and final
temperatures was taken as the temperature of the reaction.
The experiment was repeated for five different temperatures, keeping all other
variables constant.
The results were tabulated and from the results a graph of 1000/ time (y-axis)
against temperature (x-axis) was plotted. (1000/ time is a measure of the rate)
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Results:
Initial temp.
/oC
Final temp.
/oC
Mean temp.
/oC
1000/time /s-
Time /s
1
Interpretation of Results:
The graph is an exponential curve: the rate of reaction increases rapidly with
temperature, a rise in temperature of 10oC approximately doubling the rate.
The higher the temperature, the faster the reaction
In the reaction between thiosulphate
ions and hydrogen ions, as the ions
collide covalent bonds are broken
and new bonds are formed. Since
energy is needed to break bonds,
the colliding particles must have a
minimum
energy
on
collision
sufficient to break the bonds. This
energy is known as the activation
energy (Ea).
bonds
breaking
bonds
forming
Ea
Chemical
Energy
reactants
H
products
Only those collisions with energy
greater than or equal to the
activation energy will result in a
reaction.
Reaction path
An increase in temperature increases the rate of reaction in two ways:
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The particles collide more energetically
The particles of a substance have a range of different energies, the average
energy being proportional to the temperature in Kelvins. As the temperature
increases, the particles move faster (i.e. they have more kinetic energy), and the
proportion of particles with higher energies increases. Therefore, the number of
collisions with energy greater than or equal to the activation energy rises rapidly
as the temperature increases, and so the rate rises rapidly. This is the major
effect.
The particles collide more frequently
The particles move around faster and therefore there is a greater chance that
they will be involved in a collision. This is a minor effect.
3. The Concentration of Reactants
The reaction used to study the effect of concentration is the reaction of sodium
thiosulphate solution with dilute hydrochloric acid.
Na2S2O3(aq) + 2HCl(aq)
2NaCl(aq) + S(s) + H2O(l) + SO2(g)
The reaction is followed by the appearance of colloidal sulphur as the reaction
proceeds. The formation of sulphur begins as soon as the reactants are mixed, but
it takes time for observable amounts to be produced. The time taken to reach a
particular point in the reaction can be determined by standing the reaction flask on
a piece of paper marked with a feint cross and timing how long it takes for the
cross to be obscured when looked at from above.
Method:
50cm3 of sodium thiosulphate stock solution (concentration 40g per litre) were
measured out into a measuring cylinder and poured into a 250cm 3 conical flask. The
flask was placed on a piece of paper marked with a feint cross. 5cm 3 of 2M
hydrochloric acid were measured out into a second measuring cylinder. As the acid
was poured into the conical flask, a stopwatch was started and the mixture gently
swirled. The stopwatch was stopped when the cross, viewed from above, became
obscured. The time was recorded.
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Volume of Na2S2O3
(40g/l) /cm3
50
40
30
20
10
Volume of H2O
/cm3
0
Time /s
1000/ time
/s-1
10
20
30
40
The experiment was repeated for five different concentrations, keeping all other
variables constant.
A graph of 1000 /time (y-axis) against volume of Na2S2O3 (x-axis) was plotted.
Since the total volume of the reaction mixture is constant (at 55cm 3) the
concentration of Na2S2O3 is proportional to its volume.
Results:
Constants:
total volume of Na2S2O3 solution
2M hydrochloric acid
temperature
conical flask & cross
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50cm3
5cm3
20oC
x
30
x
20
x
1000/ time /s-1
x
10
x
0
10
20
30
Volume of Na2S2O3 /cm3
40
50
Interpretation of Results:
Your graph should be a straight line through the origin, therefore the rate of the
reaction is directly proportional to the concentration of sodium thiosulphate.
The more concentrated the solution, the faster the reaction
To react, the reacting particles must collide; therefore the rate will be faster the
greater the number of collisions there are in a given volume in a given time. The
more concentrated the solution is, the greater the number of particles there are in
a given volume and therefore the greater the frequency of collisions. It is
important to note that only a very small proportion of the total number of collisions
is successful and leads to a reaction.
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Effect of Pressure
Pressure is important only in reactions involving gases. Pressure affects gaseous
reactions in the same way that the concentration affects reactions in solution. As
the pressure is increased, the greater the number of particles there are in a given
volume and therefore the greater the number off collisions in a given time.
Therefore, as the pressure increases, the rate increases.
4. Addition of a Catalyst
The reaction used to study the effect of a catalyst is the decomposition of
hydrogen peroxide:
2H2O2
2H2O + O2
The reaction is catalysed by several metal oxides; the compound used here is
manganese(IV) oxide. The reaction is followed by collecting the oxygen given off
and measuring its volume at regular intervals of time. The gas may be collected
either in a gas syringe or over water in a burette.
A graph of volume of oxygen (y-axis) against time (x-axis) is plotted.
A catalyst is a substance which increases the rate of a chemical reaction but is
not used up in the reaction.
A catalyst works by providing an alternative reaction route which has a lower
activation energy.
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Interpretation of Results:
The reaction is faster the more catalyst there is present; this is shown by the
increasing steepness of the curves.
The reaction takes place on the surface of the catalyst. Increasing the mass of
catalyst increases the surface area and therefore speeds up the rate.
A catalyst does not affect the outcome
of a reaction; the same product is
formed but in a shorter time. A catalyst
works by weakening bonds, which
lowers the activation energy for the
reaction. If the activation energy is
lowered, more particles have enough
energy to react and therefore the
reaction goes faster.
Ea
Chemical
Energy
Ea cat
reactants
H
products
A catalyst is not used up in the reaction
but is recovered unchanged at the end.
Reaction path
A catalyst can be used over and over again to speed up the conversion of reactants
to products. Different reactions need different catalysts.
Biological catalysts
Enzymes are large molecules that speed up the chemical reactions inside cells.
Enzymes are a type of protein and, like all proteins, they are made from long chains
of different amino acids. They catalyse reactions such as fermentation, respiration
and photosynthesis and also speed up the breakdown of fats and other food stains
in biological washing powders.
Each enzyme will only speed up one reaction as the shape of the enzyme molecule
needs to match the shape of the molecule it reacts with (the substrate molecule).
The part of the enzyme molecule that matches the substrate is called the active
site.
If the shape of the enzyme changes, it’s active site may no longer work. We say
the enzyme has been denatured. They can be denatured by high temperatures or
extremes of pH. Note that it is wrong to say the enzyme has been killed. Although
enzymes are made by living things, they are proteins, and not alive.
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Photochemical reactions
Photochemical reactions get the energy the need from light. The more intense the
light, the faster the reaction goes. Two examples of photochemical reactions are
photosynthesis and the reaction in film photography.
Photosynthesis turns carbon dioxide and water into sugar and oxygen. It is
catalysed by chlorophyll. You have probably done an experiment in Biology lessons
to show the effect of changing light intensity on the rate of this reaction.
In photography, light falls on a film coated with silver bromide. When light is
exposed to the film, the silver ions break down (reduced) to silver metal. The more
light that hits the film, the more silver ions are reduced to silver. When the film is
developed any unreacted silver bromide is removed leaving silver behind on the
film. So the areas exposed to most light have the most silver and appear darkest.
This is what we call the ‘negative’.
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Topic 10: The speed of a reaction
Summary questions
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