Factors that Effect Rate
Now that we know what has to happen for a reaction to take place (a collision between molecules with the correct orientation and minimum activation energy), let’s examine what factors
we can manipulate in order to change the rate of the reaction.
Temperature
We’ll start by looking at a simple reaction; the dissolving of an Alka-Seltzer tablet in water.
When you ―plop‖ the tablet in water, the water molecules dissolve the citric acid and sodium
bicarbonate in the tablet and they react to form sodium citrate, water, and carbon dioxide gas
which causes the ―fizz‖. We need a baseline experiment so we’ll fill a beaker with 500 mL of
tap water and put in an Alka-Seltzer tablet and see how long it takes for it to dissolve.
Citric acid and sodium bicarbonate react to form carbon dioxide
gas which causes the ―fizz‖ effect.
HC6H7O7 + NaHCO3 —> NaC6H7O7 + H2O + CO2 (g)
This can only happen in an aqueous environment, though, so as
soon as the water dissolves the tablet, the citric acid and sodium
bicarbonate are allowed to react. The faster the water can dissolve the tablet, the faster the reaction goes.
http://marketplace.publicradio.org/display/web/2008/06/27/
financial_hangover/
Temperature (oC)
Time (seconds)
23
49
49 seconds. Great. Now let’s repeat the experiment except this time we’ll use 500 mL of boiling hot water.
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In this experiment we are using 500 mL of boiling hot water. The time is reduced to 20 seconds.
Temperature (oC)
Time (seconds)
100
20
One last experiment by using the same amount of water that is ice-cold. The time jumps up to
84 seconds.
Temperature (oC)
Time (seconds)
2
84
To summarize: hotter temperatures make reactions take place faster. Colder temperatures
make reactions take place slower. But why?
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http://www.astronomynotes.com/solarsys/s3.htm
First of all, hotter temperatures make the molecules move faster. The faster the molecules
move, the greater the number of collisions will take place. The greater the number of collisions
that occur, the greater the chance that the molecules will collide with the correct orientation and
thus increase their chances of having the collision be effective.
More importantly, though, is the fact that a higher temperature means the molecules will be colliding with each other with more energy because the kinetic energy or speed of the molecules is
directly proportional to the temperature. Consequently, a greater number molecules will have
the minimum activation energy to make the reaction proceed. Let’s look at a graph:
Number of molecules with a given energy
Ea, activation energy
300 K
In this example, we see that
none of the molecules have
enough energy to get past the
activation energy barrier, Ea.
Thus the reaction is NOT proceeding.
Increasing Energy
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Number of molecules with a given energy
Ea, activation energy
300 K
By increasing the temperature, the distribution of molecular speeds widens. Fewer
molecules are at any one
speed but there is a wider
range of speeds. Consequently, some of the molecules have enough energy to
get over the activation energy
barrier and thus the reaction
occurs slowly.
500 K
Increasing Energy
Number of molecules with a given energy
Ea, activation energy
300 K
500 K
700 K
Increasing the temperature
even more allows for more
molecules to have the minimum activation energy necessary so the reaction proceeds
even faster.
Increasing Energy
A faster temperature means the molecules move
faster. Thus, there are more collisions and thus
a greater chance for a collision with the correct
orientation.
At the higher temperature, more molecules are
past the activation energy line. The reaction
proceeds faster.
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Surface Area
Let’s stay with the Alka Seltzer experiment for just a minute more. In addition to changing the
temperature of the system, we could also grind up the tablet in a mortar. If we compare the
time it takes for the ground up tablet to the whole tablet we see a big difference in dissolving
time. The ground up tablet reacted a lot faster. Why?
Tablet
Temperature (oC)
Time (seconds)
Whole
23
49
Ground up
23
20
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Remember that in order for a reaction to proceed, the particles have to collide with each other.
When the table is whole, the particles on the inside of the tablet cannot collide with the water
molecules until the particles on the outside of the tablet have reacted away. They have to wait.
Although they have enough energy to react, they cannot because no contact can be made.
When the tablet is ground up, it exposes more surface area of the reactive molecules. This allows for more molecules to collide with each other at the same time as no molecules have to
wait until they are exposed. Thus, the reaction proceeds much faster.
Water molecules can hit the outer
layer of the tablet...
… but not these here in the center
http://www.chemguide.co.uk/physical/basicrates/surfacearea.html
But now that the tablet has been
ground up, the water molecules can
react with more of the molecules in
the tablet at once. Thus, the reaction
is faster.
http://dustexplosions.blogspot.com/2008/12/house-keeping-facility-combustible-dust.html
More surface area means more collisions can take place at the same time. More collisions
means there is a greater chance for an effective collision so the rate increases.
Every year there are industrial explosions at factories,
farms, and other places that generate a lot of ―dust‖. This
could be from something simple like flour particles to simple dust kicking up from falling grain, corn, etc. All the
exposed surface area from the minute, little pieces can turn
a nearby tiny spark into a huge explosion!
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Concentration
What would happen if I put a piece of magnesium ribbon into a beaker of hydrochloric acid?
The magnesium would react to form magnesium chloride and hydrogen gas:
Mg (s) + 2 HCl (aq) —> MgCl2 (aq) + H2 (g)
Let’s do an experiment where we take a piece of magnesium ribbon and put it in two beakers
that have the same amount of HCl at the same temperature but have different concentrations.
We’ll time to see how long it takes for the Mg to dissolve. In the first beaker we have 1 M HCl
and in the second beaker we have 6 M HCl. As we can see from the results, it took a lot longer
for the 1 M HCl to react away the magnesium. Why?
Concentration (M)
Time to Dissolve Mg Piece
(seconds)
1
344
6
14
As we have stated before, in order for a reaction to proceed, the particles must collide. The
more collisions there are, the better chance there is to have an effective collision and so the reactions will speed up. As we can see in the diagrams below, there are more particles in the concentrated solution than in the dilute one. When the magnesium is put into the beakers, there is a
much greater chance of a collision occurring in the concentrated one. Thus, the reaction goes
faster in the concentrated solution.
Dilute
Concentrated
A greater concentration means there is a greater chance of a collision. The greater the
chance of a collision means the reaction will proceed faster.
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Catalysts
Let’s examine a car trip that a salesman has to make every week. He starts in his home city and
has to drive over a big hill to get to his destination to sell his products. He knows it is going to
be profitable but, still, he has to drive his car over that great big hill all the time.
Start
Finish
What can he do about it? By himself, nothing. But if he can exert enough influence, maybe he
can get somebody to build a tunnel through the hill so nobody has to go over it.
Now, it requires much less energy for the
car to go from start to finish and it will take
the salesman less gas and time to make the
trip!
Start
Finish
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The previous page was an analogy of a catalyst– a substance that speeds up the rate of a reaction but is not ultimately consumed by the reaction. A catalyst provides the reaction an alternate pathway which lowers the activation energy, making it easier for the reaction to proceed.
We can see these effects on the two graphs we have looked at thus far during the unit.
The activation energy is
much higher for the uncatalyzed pathway here….
…. than for the catalyzed
pathway here.
On this graph we see that
the original activation energy is here...
…. but by using a catalyst
the activation energy drops
to this spot, allowing far
more molecules to have the
activation energy and thus
the reaction goes faster.
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Remember that catalysts participate in but are not ultimately consumed by the reaction. What
that means is that catalysts are often used up in one step and then re-produced in another step.
The decomposition of ozone by chlorofluorocarbons (CFCs) is a perfect example of this.
We see that Cl is acting as a catalyst. It is a reactant here in step 1
Overall:
O3
+
O
—>
2 O2
Step 1:
Cl
+
O3
—>
ClO
+
O2
Step 2:
ClO
+
O
—>
Cl
+
O2
And it is a product here in step 2.
It does not appear in the overall
equation.
Step 1:
Cl
+
O3
—>
ClO
+
O2
Step 2:
ClO
+
O
—>
Cl
+
O2
When the 2 steps are added together, the Cl drops out of the equation.
It is also good to note here that the opposite of a catalyst is also present in this reaction. This is an intermediate– a
substance that is produced in one step and then re-used in a subsequent step. Thus neither intermediates or catalysts appear in the overall equation.
ClO is an intermediate. It is made
here in step 1...
Step 1:
Cl
+
O3
—>
ClO
+
O2
Step 2:
ClO
+
O
—>
Cl
+
O2
… and re-used here in step 2. It does
not appear in the overall equation.
Step 1:
Cl
+
O3
—>
ClO
+
O2
Step 2:
ClO
+
O
—>
Cl
+
O2
Combined:
Cl + ClO + O3 + O —> Cl + ClO + 2 O2
Overall:
O3
+
O
—>
2 O2
Note the catalyst in red and
the intermediate in yellow
drop out.
A catalyst lowers the activation energy by providing a different pathway for the reaction. The
reaction speeds up because more molecules have the minimum activation energy. The catalyst
is consumed and then re-produced in the reaction so does not appear in the overall equation.
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