HPAVC/HPAVC CD-ROM.iso/pc/HOMEWORK.ZIP/381.TXT

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Essay Name       : 381.txt
Uploader         :
Email Address    : [e-mail]
Language         : english
Subject          : Chemistry
Title            : rates of reaction
Grade            : 90%
School System    :
Country          : uk
Author Comments  :  An excellent project, very perceptive
Teacher Comments :
Date             : 29/11/96
Site found at    : friend
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  BACKGROUND INFORMATION


 What affects the rate of reaction?
1)  The surface area of the magnesium.
2)  The temperature of the reaction.
3)  Concentration of the hydrochloric acid.
4)  Presence of a catalyst.

 In the experiment we use hydrochloric acid which reacts with the
magnesium to form magnesium chloride.  The hydrogen ions give
hydrochloric acid its acidic properties, so that all solutions of
hydrogen chloride and water have a sour taste; corrode active metals,
forming metal chlorides and hydrogen; turn litmus red; neutralise
alkalis; and react with salts of weak acids, forming chlorides and the
weak acids.

 Magnesium, symbol Mg, silvery white metallic element that is
relatively unreactive. In group 2 (or IIa) of the periodic table,
magnesium is one of the alkaline earth metals. The atomic number of
magnesium is 12.

Magnesium(s) + Hydrochloric acid(aq) = Magnesium Chloride(aq) +
Hydrogen(g)
  Mg     +    2HCl        =     MgCl2
      +     H2

 In the reaction when the magnesium hits the acid when dropped in,
it fisses and then disappears giving of hydrogen as it fisses and it
leaves behind a solution of hydrogen chloride.
 The activation energy of a particle is increased with heat.  The
particles which have to have the  activation energy are those particles
which are moving, in the case of magnesium and hydrochloric acid, it is
the hydrochloric acid particles which have to have the activation energy
because they are the ones that are moving and bombarding the magnesium
particles to produce magnesium chloride.

 The rate at which all reactions happen are different.  An example
of a fast reaction is an explosion, and an example of a slow reaction is
rusting.
 In any reaction,
   reactants chemical reactions« products.

 We can measure reactions in two ways:
1)  Continuous:- Start the experiment and watch it happen; you can use a
computer ôloggingö system to monitor it.  I.e.  Watching a colour fade or
increase.
2)  Discontinuous:-  Do the experiments and take readings/ samples from
the experiment at different times, then analyse the readings/samples to
see how many reactants and products are used up/ produced.

 Reaction rate = amount of reactant used up
    time taken

 If the amount used up is the same each time then the only thing
that changes is the time taken.
  so,  reaction rate ╡     1
          time taken.
           rate =       K
           time taken.
 Where K is the constant for the reaction.

 For particles to react:-
a)  They have to collide with each other.
b)  They need a certain amount of energy to break down the bonds of the
particles and form new ones.  This energy is called the ôActivation
Energyö or Ea.
 When we increase the temperature we give the particles more
energy which:
1)  Makes them move faster which In turn makes them collide with each
other more often.
2)  Increases the average amount of energy particles have so more
particles have the ôactivation energyö
 Both of these changes make the rate of reaction go up so we see a
decrease in the amount of time taken for the reaction and an increase in
 1

time taken.
  =     1
         time taken.     Reflects the rate of reaction.
 Because temperature has an effect on both the speeds at which the
particles react and the activation energy they  have a greater effect on
the rate of reaction than other changes.

 A change in concentration is a change in the number of particles
in a given volume.
 If we increase the volume:-
a)  The particles are more crowded so they collide more often.
b)  Although the average amount of energy possessed by a particle does
not change, there are more particles with each amount of energy;- more
particles with the activation energy.
 a) is a major effect which effects the rate, but b) is a minor
effect which effects the rate very slightly.
  In this experiment we are not concerned with whether the reaction
is exothermic or endothermic because we are concerned with the activation
energy needed to start and continue the reaction.




    PREDICTIONS

 I predict that as we increase the temperature the rate of
reaction will increase.
 If we increase the temperature by 100C the rate of reaction will
double.

 I predict that if we increase the concentration of the acid the
reaction rate will increase.
 If the concentration of the acid doubles, the rate of the
reaction will also double.


   LINKING PREDICTION TO
     THEORY

Reaction Rate and Temperature.
 The collision theory describes how the rate of reaction increases
as the temperature increases.  This theory states that as the temperature
rises, more energy is given to the particles so their speed increases,
this increases the number of collisions per unit of time.  This increase
in collisions increases the rate of reaction.
 The collision theory explains how the rate of reaction increases,
but it does not explain by how much or by how fast the rate increases.
The Kinetic energy of a particle is proportional to its absolute (Kelvin)
temperature.
   1/2 mv    2╡ T
 But the mass of the particles remains constant so we can
eliminate that part of the equation so;
  ▐ V2╡T
 Therefore we can fit this into a formula:
  V21/V22 = T1/T2
 If we substitute the temperature into the formula we can work out
the average speed of the formula:
  V21/V22 = 310/300
  \V 1    =  ╓310/300V 2
    = ╓1.033V2
     =  1.016V2

 However if we look at this it is only 1.016 times greater than
the speed at 300K, in other words we can see that it has only increased
by 1.6%.
 The frequency of the collisions depends on the speed of the
particles, this simple collision theory only accounts for the 1.6%
increase in the rate, but in practice the reaction rate roughly doubles
in a 10K rise, so this simple theory cannot account for an 100% increase
in the reaction rate.
 During a chemical reaction the particles have to collide with
enough energy to first break the bonds and then to form the new bonds and
the rearranged electrons, so it is ôsafeö to assume that some of the
particles do not have enough energy to react when they collide.
 The minimum amount of energy that is needed  to break down the
bonds is called the activation energy (EA).  If the activation energy is
high only a small amount of particles will have enough energy to react so
the reaction rate would be very small, however, if the activation energy
is very low the number of particles with that amount of energy will be
high so the reaction rate would be higher.  An example of a low EA would
be in explosives when they need only a small input of energy to start
their exceedingly exothermic reactions.
 In gases the energy of the particles is mainly kinetic, however
in a solid of a given mass this amount of energy is determined by their
velocities.


 This graph below shows how the energies of particles are
distributed.









 This graph is basically a histogram showing the number of
particles with that amount of energy.  The area underneath the curve is
proportional to the total number of particles.  The number of particles
with > EA is proportional to the total area underneath the curve.
 The fraction of particles with > EA  is given by the ratio:
   Crosshatched area under the curve
    total area under curve

 Using the probability theory and the kinetic theory of gases,
equations were derived for the distribution of kinetic energy amongst
particles.  From these equations the fractions of particles with an
energy > EA J mole-1 is represented by the equation: e   -Ea/RT  where R=
the gas constant (8.3 J K-1 mole -1)
      T= absolute temperature.
 This suggests that at a given temperature, T,
  The reaction rate ╡ e   -Ea/RT
 If we use k as the rate constant, as a measure of the reaction
rate we can put this into the equation also.
 k╡ e   -Ea/RT
 ▐ k= A e   -Ea/RT
 The last expression is called the Arrhenius equation because it
was developed by Srante Arrhenius in 1889.  In this equation A can be
determined by the total numbers of collisions per unit time and the
orientation of the molecules when the collide, whilst e   -Ea/RT is
determined by the fraction of molecules with sufficient amounts of energy
to react.
 Putting the probability theory and the kinetic theory together
this now gives us a statement which accounts for the 100% increase in the
rate of reaction in a 10K rise.


Reaction Rate and Concentration.
 The reaction rate increases when the concentration of the acid
increases because:
If you increase the concentration of the acid you are introducing more
particles into the reaction which will in turn produce a faster reaction
because there will be more collisions between the particles which is what
increases the reaction rate.


    METHOD.

 To get the amount of magnesium and the amount of hydrochloric
acid to use in the reaction, we have to use an excess of acid so that all
of the magnesium disappears.
  Mg     +    2HCl        =     MgCl2
      +     H2
        1 mole         2 moles            1 mole
 1 mole

 So, we can say that one mole of magnesium reacts with 2 moles of
hydrochloric acid.
 If we use 1 mole of magnesium and 2 moles of hydrochloric acid we
will get a huge amount of gas, too much for us to measure.  We would get
24,000 cm3 of hydrogen produced where we only want 100 cm3 of hydrogen
produced.  So to get the formula for the amount of moles that we have to
use the formula:
 Moles = mass of sample         100    =0.004 moles.
  volume with 1 mole        24,000

To get the maximum mass we can use:
 Mass = moles x RAM.
         = 0.004 x 24
         =  0.0096g
 So, this is the maximum amount of magnesium we can use.  To the
nearest 0.01 of a gram =  0.01.  This is the maximum amount of magnesium
we can use.
 Because the reaction reacts one mole of magnesium to two moles of
hydrochloric acid we have to make sure that even with the lowest
concentration of acid we still have an excess of acid.
 The acid that we were using was 2 moles per dm2 which means that
it is 0.2 moles per 100 cm2 of acid.
 We need to make the reaction work to have double the amount of
magnesium.  The maximum number of moles that the magnesium needed was
0.004 moles so the amount of acid that we needed was double that so that
equals 0.008 moles.  As you can see from the table below we have the acid
in excess throughout the experiment.
Amount of HCl (cm3)    Amount of H2O (cm3)    Moles of acid.
100    0    0.2
75    25    0.15
50    50    0.1
25    75    0.05
The reason why we used 0.01g of magnesium was because it was therefore
easy to measure because there was not too much, or too little.  Therefore
we had no problem with too much gas.




Apparatus
  This is the apparatus we used to measure the amount of H2 that was
produced in the reactions.  We measured the amount of gas that was given
of  every two seconds to get a good set of results.  We used this
apparatus with the reaction changing the concentration, and then the
temperature. To accurately measure the amount of gas given of we used a
pen and marked on the gas syringe at the time intervals.
 This is the apparatus we used to measure how long it took for the
magnesium to totally disappear.  We used this apparatus in both of the
experiments, changing the temperature and the concentration of the acid
to water.

Temperature.
 When we did the experiment changing temperature we used both of
the sets of apparatus.  To get a fair reaction we had to keep the amount
of magnesium the same and the concentration of the acid.  In the
experiment we used 0.1g of magnesium and the concentration of the acid
was 50cm3 of acid to 50cm3 of water.  This is because if we used 100cm3
of acid the reaction would be too fast. Still we had an excess amount of
acid, so one mole of magnesium can react with two moles of HCl.
Concentration.
 When we did the reaction changing the concentration we changed
the concentration until we had just enough for 1 mole of magnesium to
react with two of HCl.  To get  a fair reaction we had to keep the amount
of magnesium the same and the temperature.  We used 0.1g of magnesium.












     RESULTS
Temperature

 From this graph you can see that if we do increase  the
temperature the rate of reaction also increases, but it does not show
that if you increase the temperature the rate of reaction doubles.

 This graph shows that there is an increase in the rate of
reaction as the temperature increases.  This shows a curve, mainly
because our results were inaccurate in a number of ways.  This is because
the concentration is changed during the experiment because at high
temperatures the acid around the magnesium is diluted.  If this
experiment was accurate it would be also a curve but if you made it into
1/time the result would be a straight line showing a clear relationship.

Even though I changed it to 1/time it still does not show a clear
relationship because of the factors mentioned in the conclusion.
Concentration

This graph shows an increase in the amount of gas given off and the speed
at which it is given off.  This graph also does not show the rate
increase, it just shows how it increases with a change in concentration.

This graph shows that if you increase the concentration of the molar
solution of the acid the time in which the Mg takes to disappear becomes
a lot slower.  This does not show the rate at which this happens, the
graph of rate vs. conc. would show a straight line.

This shows a straight line, thus proving that there is a relationship
between the time it takes the magnesium to disappear and the
concentration of the acid.  If we take a gradient of it, it would show
the rate at which the reaction was happening.

 Because this shows a straight line we can say that it is a second
order reaction.

 This graph shows a nearly straight line which shows that there is
a relationship between  the temperature  and the rate of reaction,  as
the gradient shows the rate of reaction.  If you look at this graph it
comes out to show that if you increase the temperature by 100C the
gradient of the line is doubled.  This shows that rate ╡ temp.

This graph shows that if you increase the molar concentration of the
acid, you will increase the rate of reaction.  From this you can see from
the gradient, that if you double the molar concentration of the acid the
rate of reaction will double because the gradient is a way of showing the
rate of reaction.
 If you compare the quantitative observations to see which the
faster reaction is you can see that after 10 seconds:
Temp.    2    10    20    30    40    50
Amount of H2 produced after 10s    7.5    16    25    54    57
83
 Even though there is a greater increase in the amount of H2 given
off in each of the different reactions you can see that there is a change
in the amount given off, but between the temperatures 30 and 400C there
is not much of a change, this could be because of our human error, there
should be a big change in the amount given off.
Molar conc.    0.5    1    1.5    2
Amount of H2 produced after 10s    6    25    60    90
This table shows a nice spread of results throughout the range of
concentration.  It clearly shows that the reaction is at different stages
so is therefore producing different amounts of H2.  This shows also that
the reaction is affected by the concentration of the acid.


           CONCLUSIONS

 I conclude that if you increase the temperature by 10oC the rate
of reaction would double, this is because of using the kinetic theory and
the probability theory.  Even though our results did not accurately prove
this, the theory that backs it up is sufficient.  the kinetic theory
explains that if you provide the particles with a greater amount of
kinetic energy they will collide more often, therefore there will be a
greater amount of collisions per unit time.  The probability theory
explains that there is only a number of particles within the reaction
with the amount of Ea to react, so if you increase the amount of kinetic
energy there will be more particles with that amount of Ea to react, so
this will also increase the reaction rate.
 If you double the concentration of the acid the reaction rate
would also double, this is because there are more particles in the
solution which would increase the likelihood that they would hit the
magnesium so the reaction rate would increase.  The graph gives us a good
device to prove that if you double the concentration the rate would also
double.  If you increase the number of particles in the solution it is
more likely that they will collide more often.
 There should be more H2 given off if we compare it across the
range of temperatures because the reaction is going quicker and so more
H2 is given off in that amount of time.
 There is more H2 given off if you compare it to the range of
concentrations that you are using, this shows that the reaction is at
different stages and so is therefore producing different amounts of H2.
 Also our results were not accurate but this could be because of a number
of reasons.
 There our many reasons why our results did not prove this point
accurately.
╖ At high temperatures the acid around the magnesium starts to starts to
dilute quickly, so if you do not swirl the reaction the magnesium would
be reacting with the acid at a lower concentration which would alter the
results.
╖ Heating the acid might allow H Cl to be given off, therefore also
making the acid more dilute which would also affect the results.
╖ When the reaction takes place bubbles of H2 are given off which might
stay around the magnesium which therefore reduces the surface area of the
magnesium and so the acid can not react properly with it so this affects
the results.

 To get more accurate results, we could have heated the acid to a
lower temperature to stop a large amount of H Cl being given off.  The
other main thing that could have helped us to get more accurate results
is we cold have swirled the reaction throughout it to stop the diluting
of the acid and the bubbles of H2 being given off.

 If I had time I could have done the reactions a few more times to
get a better set of results.  This would have helped my graphs to show
better readings.

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