Independent variable: Concentration of Hydrochloric acid, temperature at which reaction takes place and surface area.-changing one at a time in order to attain reliable results.
Dependent variable: Duration of reaction, and therefore the rate of reaction.
The complexity of our variables are as follows:
Temperature shall be a continuous variable as we shall begin with 220c (room temperature). Concentration shall also be a continuous variable as we shall begin with ½ mol/litre solution. The affect of surface area shall. also be continuous as we shall be calculating the surface area of our reactants rather than only using the categoric variables of looking at powder and tablet form.
Predictions:
In order to predict the outcome of our investigation it is necessary to consider how and why chemical reactions occur. According to the hard-sphere collision theory, molecules are assumed to be hard spheres and must collide in order for a reaction to occur. Yet the kinetic theory of gases shows that the frequency of collisions is so high that reactions would be virtually instantaneous if every collision led to a reaction. Svante Arrhenius therefore put forward the suggestion that only molecules which possess more that a critical energy known as the activation energy, Ea , are able to react. The activation energy enables the chemical bonds to stretch and break and rearrangements of atoms, ions and electrons to occur. The energy profile of an exothermic reaction, according to this, is as shown:
Following the progression of the above energy profile from right to left, as the molecules are approaching, there is little change in the total temperature until they get closer, hence the flat part on the left. According to the transition state theory, as the molecules gain more energy and are with a few nanometers of each other, electrostatic repulsion between their nuclei and the negative electron clouds begin to operate. In order to overcome this, the molecules must possess sufficient speed and kinetic energy. The peak, where this has been achieved, is the energy barrier which must be surmounted in order for the reaction to take place. Arrhenius based his quantitative analysis on some of these principles and hence came up with the Arrhenius equation. In most cases it is shown to be highly accurate and indeed it has a surprisingly wide applicability. For example, the law is obeyed by the chirping of crickets, the creeping of an ant, the rate of ageing and many other examples, the reason being that they are all controlled by chemical reactions. However the Arrhenius equation was shown to have many discrepancies. Thus Henry Errying came up with a more sophisticated theory known as the transition-state theory, which takes in to account other things such as molecular orientation. The peak of the energy profile, where the reactants have a high energy is described as the transition-state or activated complex, designated by the symbol . Errying introduced the idea of reaction geometry and incorporated this in his more robust equation.
Therefore, three main conditions must be met if a reaction is to occur. Firstly, the molecules must collide. Secondly, they must be positioned so that the groups are together in a transition state between reactants and products and finally, the collisions must have enough energy to form this transition state and thus convert it into products. A fast reaction therefore occurs when all three criteria are met easily. However, if even one is difficult, the reaction shall be typically slow.
Having considered the conditions necessary for a reaction we may now make predictions regarding the effect of our three independent variables.
Concentration Of Hydrochloric Acid:
I believe that an increase in reactant concentration shall lead to an increase in the rate of reaction. This is because a higher concentration means a larger number of reactant particles are present in a given volume. Thus there shall be more collisions resulting in a greater number of collisions of the correct orientation, therefore increasing the amount of successful collisions in a given time. Thus the reaction rate will increase. And yet we find that through experimentation this is not always the case. The relationship between the concentration of an individual reactant and the rate of the reaction is summed up by the rate equation or rate law:
Rate = k (X)n
Where k is a constant, known as the rate constant for the reaction, X is the reactant under consideration (in our case it is hydrochloric acid) and n is known as the order of the reaction. When n=1 the reaction is therefore directly proportional to the concentration of X, i.e. rate α (X). This is known as a reaction of the first order with respect to X. If n=2, the reaction rate is therefore proportional to X2, i.e. rate α (X)2. This is known as a reaction of the second order with respect to X. Another possibility is that n=0, therefore:
Rate = k (X)0
but (X)0 = 1
Rate = k
From this we may say that if the reaction is of zero order with respect to X, the reaction rate is independent of the concentration of X. The variations of reaction rate with concentration for reactions which are zero, first and second order is illustrated below:
Collisions in which three or more molecules all come together at the same time are very unlikely; the reaction will instead proceed more rapidly by a composite mechanism involving two or more elementary processes, each of which is only first or second order. Thus there are a few reactions of third order and reactions of a higher order than that are unknown.
Yet the order of a reaction is found to be unrelated to the stoichiometric equation and can not be worked out in anyway other than by the use of experimentation. We may say that the reaction is unlikely to be of third order, yet we are still left with three predictions, however the first is most likely as it is the most common:
1 The reaction is of first order with respect to the hydrochloric acid, thus the rate shall increase proportionally as does the concentration.
2 The reaction is of the second order with respect to the hydrochloric acid, thus the rate shall increase proportionally as the square of the concentration does.
3 The reaction is of zero order with respect to the hydrochloric acid, thus the concentration of hydrochloric acid will have no effect on the rate of reaction.
Temperature Of Hydrochloric Acid:
With regards to temperature, the kinetic theory shows us that the kinetic energies of molecules, particularly gases, cover a very wide range. Some have very small, some intermediate and few have high energies. This is explained by the random movement of particles known as Brownian motion. Thus, supposing at a given time, all the molecules in a gas had the same kinetic energy, random collisions would tend to speed up some and slow others down. Although there is a distribution of kinetic energies, at a higher temperature, given that the kinetic energy of a particle is proportional to its absolute temperaturei.e. ½ mV2 T, more particles will have energies greater than the activation energy of the reaction. Therefore there shall be a greater number of successful collisions, resulting in an increased rate of reaction. This is illustrated in the Maxwell-Boltzmann distribution of kinetic energies at two different temperatures, where Ea is the activation energy:
Temperature T2 • temperature T1 (in K)
Relative
probability of
kinetic energy
occurring Kinetic Energy
The above can be expressed mathematically:
n=noe-E/RT
where n is the number of molecules with energies greater the E, the activation energy, no is the total number of molecules, R is the gas constant, equal to 8.3145 J K-1 mol-1 and T is the absolute temperature. Although from this we may predict that increasing the temperature of the hydrochloric acid may increase the rate of reaction, it may also be necessary to quantify this prediction.
It has been found through experimentation that the rate of many reactions doubles for a temperature rise of only 10K. This would lead us to expect results similar to those below:
Reaction
Rate
Temperature
Surface Area Of Calcium Carbonate:
From the hard-sphere collision theory one may predict that the greater the total area exposed to reaction is the faster the rate of reaction shall be. This is as collisions ‘only’ occur on the surface. Providing a larger surface area will thus accommodate a greater number of collisions, thus resulting in a greater chance of successful collisions occurring, hence increasing the rate. In the case of homogenous systems, the idea of surface area becomes somewhat meaningless as reacting substances normally occur in their maximum state of subdivision. However, in the case of a heterogeneous system, such as the reaction we are investigating, the area of contact between systems will influence the rate of reaction considerably, thus there is likely to be considerable change in the rate of reaction particularly between the reactions in which the calcium carbonate is in powder or tablet form.
Several leading chemists have developed a quantitative approach towards the influence of collision cross section or surface area. Among these are the works of the German chemist Max Trautz (1880-1961) and the British chemist William Cudmore McCullagh Lewis. Both of whom, independently came up with a formula, in 1916 and 1918 respectively, concerning this topic. However, their equation may be of little use to us as there is another factor that must be considered; the rate of diffusion. In this experiment the collisions will not only be taking place on the surface of the calcium carbonate but also inside it by methods of diffusion. Yet as the rate of diffusion is dependent on the distance that has to be travelled the rate of reaction will be faster nearer the surface, obviously being fastest on the surface as no diffusion is necessary. This can be illustrated as follows:
(Diagram not to scale) The small circles within the
tablet represent small grams
which have been joined to
make the tablets,(not all are
shown). The arrows represent colliding reactant
particles.
Thus it can be seen that surface area of the tablet will have an effect on three rates(directly or indirectly), all of which will effect the rate of reaction;
1 The rate of reaction at surface of gram
2 The rate of diffusion into gram
3 The rate of diffusion into the tablet.
Thus we may predict that increasing surface area will increase the rate of reaction, however as there are several factors that will be altered quantifying this prediction presents some difficulty.
Apparatus:
- Safety goggles and plastic gloves
- Test-tubes
- Top-pan balance
- Measuring cylinder
- Scalpel
- Test-tube rack
- Stop-clock
- Bunsen burner, tripod and gauze
- Displacement can
- Funnel
- Spatula
Plan of investigation:
Before we begin the investigation it is essential that we take various safety precautions. These shall include wearing safety goggles for eye protection and wearing plastic gloves to protect our hands as the hydrochloric acid is corrosive. It is also crucial that we keep our area used for experimentation tidy to prevent accidents. It is important to keep the reactants separate whilst setting up the apparatus so that the starting time of the reaction can be measured accurately.
The investigation shall proceed in three main parts, each dealing with one different independent or input variable. We shall first consider that of concentration.
It is likely that we will have enough time to carry out the experiments to determine the effect of hydrochloric acid on the rate using five different concentrations. With the results of five different concentrations, I believe our data will be sufficient to evaluate our predictions. However our data would be more accurate if we used the method of averaging, i.e. repeat each experiment several times which is useful in order to reduce experimental error. However, it is unlikely that we will have sufficient time to use the method. In order to receive a fair set of results regarding the effect of concentration the following variables shall be constant:
1 Amount of Hydrochloric acid -15cm3
2 Form of calcium carbonate, i.e. solid or powder-(powder)
3 Temperature - room temperature
4 Mass of calcium carbonate (0.25 grams)
The steps taken shall be as follows:
Step1) We shall measure 15cm3 of ½ molar hydrochloric acid. This shall be done, as in all the experiments by pouring the hydrochloric acid into a measuring cylinder through a funnel as shown below:
It is important that we take our reading from the surface of the liquid (meniscus) to reduce error. A measuring cylinder is accurate to the nearest 1cm3. Therefore our measurement shall be 15cm3 + 1.
We may also find out the percentage error using the following equation:
Degree of accuracy x 100
Total measurement
1 x 100 = 6.7% (to 2.s.f)
15
Thus we may expect an error of approximately up to 6.7%. Although this does seem rather high, the only other measuring cylinder that we have, which is accurate to 0.1cm3, only has a capacity of 10cm3, and therefore is of no use.
Step2) We shall measure out 0.25grams of calcium carbonate. This can be done using a top pan balance, spatula and a petri dish as shown below:
It may be worth noting that the petri dish is to be placed on the top-pan balance first, at this point the balance must be ‘tarred’ or reset to 0.00 until the calcium carbonate is added using the spatula. The mass must be read once the balance has steadied. The top-pan balance is accurate to two decimal places, i.e. to one hundredth of a gram, thus our result is likely to be in the region of 0.25g + 0.01. Hence we may work out the percentage error:
0.01 x 100 = 4%
0.25
Therefore we can expect an error of approximately 4%.
Step3) We transfer the calcium carbonate from the petri-dish to a test-tube using a funnel. We then place the test-tube in a test-tube rack.
Step4) The next step is to pour the hydrochloric acid from the measuring cylinder into the test-tube. This too can be done using a funnel. The moment the hydrochloric acid is added to the calcium carbonate the stop-clock shall be started. This shall be done as below:
Step5) When the calcium carbonate dissolves and all the effervescence has ceased the stop-clock shall be stopped.
This shall be repeated four more times using 1,2,3 and 5 molar solutions of hydrochloric acid. As there is no 3 molar solution of hydrochloric acid available in our laboratory we shall have to make up a 3molar solution ourselves. This can be done by diluting a 5 molar solution using a ration of 3acid:2water, i.e.9cm2 acid and 4 of water
Our results for this experiment are to be plotted on a table with values for the concentration, the time and the rate of reaction. We shall also plot a graph of concentration against time as well as a graph of rate of reaction against concentration. Illustrating our results graphically will help us in evaluating our predictions.
Having considered concentration we may consider the affect of temperature. In this experiment the following variables shall be kept constant:
1 Amount of Hydrochloric acid -15cm3 of ½ mol/litre (this is so that
our reaction will be slow enough to be analysed)
2 Form of calcium carbonate, i.e. solid or powder-(powder)
3 Concentration of hydrochloric acid.
4 Mass of calcium carbonate (0.25 grams)
We shall vary the temperature taking five different readings at 22, 30, 40, 50, and 60 degrees Celsius. Again I believe that the data received will be sufficient to evaluate our predictions concerning the affect of temperature.
The steps taken shall be as follows:
Step1) We shall measure 15cm3 of 1 molar hydrochloric acid. This shall be done as explained earlier.
Step2) We shall measure out 0.25grams of calcium carbonate. This can be done using a top pan balance, spatula and a petri dish as illustrated earlier.
Step3) We transfer the calcium carbonate from the petri-dish to a test-tube using a funnel. We then place the test-tube in a test-tube rack.
Step4) We then pour some water into a beaker. Too much water will result in too much time being spent heating it and too little water will result in the heat being lost rapidly. Thus we shall fill up the beaker to approximately half-way. The test tube containing hydrochloric acid shall be placed inside the beaker. This is so that the hydrochloric acid accurately reaches the temperature required-as it is surrounded by water all parts of the hydrochloric acid will be the same. The water also prevents the hydrochloric acid from cooling down rapidly after the Bunsen burner has been removed.
Step5) The beaker shall be placed on a tripod and gauze so that it may be heated by a Bunsen burner. A thermometer will be used to check that the right temperature has been reached (no heating is necessary when conducting the experiment a t room temperature). The temperature should be noted when the mercury in the thermometer has steadied after the Bunsen burner has been removed. The readings taken from the thermometer provide us with an additional error factor that must be considered. The thermometer is accurate to the nearest degree. According to the working out illustrated below the percentage error encountered when measuring the temperature, for our specific experiment, lies between 1.7% and 4.5%.
1 x 100 = 4.5% 1 x 100 = 1.7%
22 60
The apparatus shall be set up as follows:
Step6) When the appropriate temperature has been reached the calcium carbonate may be poured into the hydrochloric acid, at this moment the stop-clock shall be started. When the reaction ceases the watch shall be stopped.
The steps described above shall be carried out five times at temperatures of 22, 30, 40, 50 and 60 degrees Celsius.
The results for this experiment may be plotted on a table with values for the temperature, the time and the rate of reaction. We shall also plot a graph of temperature against time and one of rate of reaction against temperature.
Finally we must consider the variable of surface area. The method for this experiment shall be somewhat complex. We shall attempt to use five different surface areas of relatively the same mass and volume of calcium carbonate rather than only investigating the affect of having the calcium carbonate in two forms. In order to attain reliable results, as it is highly likely that there shall be an increased experimental error in this experiment, I have decided to use a mass of 0.65 grams of calcium carbonate, rather than 0.35 grams in order to decrease the percentage error and hence significance of the errors encountered in this experiment. Yet this will be at the expense of time- as with a greater mass the time for it to fully dissolve will take longer.
Thus the variables that shall be kept constant are as follows: 1 Amount of Hydrochloric acid -15cm3 2mol/litre solution.
2 Temperature
3 Concentration of hydrochloric acid-2mol/litre
4 Mass of calcium carbonate (0.65 grams)
The steps taken shall be as follows:
Step1) We shall begin by acquiring four pieces of solid calcium carbonate all weighing 0.65 grams (the percentage error has been reduced from 4% to 1.5%). The shape (roughly spherical) and hence the surface area should also be relatively the same.
Step2) Provided that the four pieces are of the same mass and hence the same volume and approximately the same surface area, we may measure the volume of one them (taken as representative of them all) by placing it into a displacement can containing acetone (this will not react with the calcium carbonate unlike water)and measure the volume of the displaced acetone. This may be done as below:
Step3) We may vary the surface area of these pieces by cutting them into several individual pieces. The greater the number of pieces the greater the total surface area. The first piece shall not be cut, the second shall be broken into two, the third into three and the fourth into four pieces-all pieces will be used, i.e. with the second piece, two pieces of calcium carbonate will be placed in the hydrochloric acid. The calcium carbonate shall be cut using a scalpel, the greater the number of pieces the greater the surface area. In order to be attain values for the surface area we must use various mathematical calculations. The pieces of calcium carbonate are closest in shape to spheres. By assuming that they are spheres we may calculate the surface are given that we have a value for the volume:
Volume of sphere = 4 πr3 ,surface area = 4 πr2
3
hence given V (volume) we must find r:
r = 3 v
4/3π
substituting into the surface area equation gives:
Surface area = 4 π x 3 v 2
4/3π
Using this equation we may obtain the first value for surface area.
Using two piece, resulting in approximately half the volume for each piece the equation shall be as follows
4 π x 3 ½v 2 x 2
4/3π
The total surface area of the third (split into 3 pieces) may be calculated using this equation:
4 π x 3 1/3v 2 x 3
4/3π
The total surface area of the fourth piece (split into four pieces) may be calculated using this equation:
4 π x 3 1/4v 2 x 4
4/3π
Step4) 0.65g of powdered calcium carbonate (prepared by the laboratory technician) may be weighed and placed whilst in a petri-dish and then placed in a test-tube using a funnel. This will give us our fifth and largest surface area.
Step5) I shall prepare 5x 15cm3 of hydrochloric acid of 2 mol/litre concentration. This concentration shall be used in order to counter the effect of using an increased mass, i.e. to save time.
Step6) The hydrochloric acid shall be added to each of the different surface areas and the time taken for the reactions to take place noted.
The results may be illustrated on a table of surface area, time and reaction rate. We may plot a graph of surface area against time and one of rate against surface area. Provided that we may draw a line of best fit with our data, it may be possible to estimate the surface area of the powdered calcium carbonate.
Observations/Problems
Each reaction performed involved fizzing due to the release of carbon dioxide (CO2) gas. It was noticed that the faster the time of reaction the quicker the bubbles emerged. The reaction was at its fastest at the start then slowed down slightly. In one or two of the experiments we encountered a slight difficulty in transferring the calcium carbonate from the petri dish to the test tube using the funnel. Also, in the experiment regarding surface area we were faced with two problems. The first was minor, we found that when cutting the Calcium carbonate into several pieces small pieces broke off as well. Although the pieces were of almost negligible size, they may add to our experimental error. The second problem was finding the volume of the calcium carbonate using the displacement can. We found that it was very hard to fill the displacement can with the right amount of acetone and thus ensure that the acetone would come out of the can when the calcium carbonate was added. We overcame this problem by using a measuring cylinder of small capacity but increased accuracy rather than using the displacement can. The volume of the calcium carbonate came out to be
0.2cm3. Using the formulas mentioned earlier, the total surface area of calcium carbonate used in each experiment is:
One piece: 4.6cm2
Two pieces: 7.2 cm2
Three pieces: 9.5 cm2
Four pieces 11.5 cm2
Results
Our results may be tabulated as follows:
Input Variable 1: Concentration Input Variable 2: Temperature
Concentration Time Rate Temperature Time Rate
mol/litre s s-1 0C s s-1
0.5 167 0.006 22 170 0.0069
1 77 0.013 30 97 0.010
2 42 0.024 40 58 0.017
3 27.8 0.036 50 22 0.045
5 14.3 0.070 60 14 0.071
Input Variable 3: Surface Area
Surface area Time Rate
cm2 s s-1
4.6 934 0.0011
7.2 780 0.0013
9.5 720 0.0014
11.5 580 0.0017
Powder 42 0.024
Our results may also be illustrated graphically:
Conclusion
The results obtained show much of what was predicted to be correct. With regards to the effect of concentration, the results clearly show that when using an increased reactant concentration, the duration of the reaction is decreased and hence the rate of reaction is increased. Thus the acid is clearly not of zero order with respect to the reaction. That is to say that in this experiment the hydrochloric acid took part in the slowest stage of the reaction mechanism known as the rate-determining step or rate-limiting step.
By analysing the graph showing rate of reaction against concentration, we may express the relationship between the concentration and the rate of reaction using the rate equation or rate law. For our reaction (since it is a neutralisation reaction):
Rate = k (CO32-)n (H+)m
We have found that m may be replaced by one. This is as the graph approximately shows linear proportionality between reaction rate and concentration. Hence, the results show that the reaction is first order with respect to the hydrochloric acid.
The results obtained for this experiment seem fairly accurate. By comparing our points with that of a trendline we find that most of the points lie within an 8% error margin. This was largely what was expected since our experiment involved several errors i.e. 6.7% in measuring the hydrochloric acid, 4% in weighing the calcium carbonate additional error may also be accounted for by the stopclock.
Regarding the effect of temperature our results also show, in agreement with our predictions, that an increase in temperature reduces the duration of the reaction and increases the rate of reaction exponentially. By using the figures in which we increased the temperature by ten degrees we may find the average (mean) increase in time and hence rate:
97 + 58 + 22
58 22 14 = 1.96
Thus when increasing the temperature by ten degrees our experiment has shown that there is a mean in rate of 96%. Bearing in mind experimental error, this agrees with my initial prediction that an increase in ten degrees doubles the rate of reaction. The experimental error in this experiment was thus rather low, a mean error of only 4%. Looking at the chart with the inserted exponential trend-line, we find that individually, each point (except perhaps one) is within the eight percent error bars.
Yet how can we explain this rapid increase in rate? Taking the equation ½ mV2 T, we may say that as the mass of a given particle remains constant
⇒ V2 T
1V2 = T1
2V2 T2
where 1v is the velocity at temperature T1, and 2v is the velocity at temperature T2. Thus by substituting in this equation, supposing that the average speed of a particle is v at 293K (20oc) we may find the percentage increase of the average speed at 303K (30oc):
1V2 = 303 1V = √ 303 = 1.017v (to 3.d.p)
2V2 293 293
From this we find that raising the temperature by ten degrees has only resulted in a 1.7% increase in average speed. On the assumption that the frequency of collisions is dependent on the average speed of the particles, similarly we may expect the rate of collisions and thus the rate of reaction to increase by 1.7% as well. Yet according to our experiment this is not so.
In order to explain this rapid increase in rate we may refer back to the equation obtained from the Maxwell-Boltzmann distribution:
n=noe-E/RT
From this equation and further work studied by Svante Arrhenius and independently by Van’t Hoff, earlier in 1884, it is possible to obtain what is known as the Arrhenius equation, bearing in mind that the rate constant is proportional to the rate of reaction:
k = Ae-E/RT
where k is the rate constant and A (the Arhenius constant) is known as the pre-exponential factor, which can also be regarded as the collision frequency and orientation factor, whilst e-E/RT represents an activation state factor.
Since many common reactions have an activation energy of 85kJ mol-1, simple substitution in the Arrhenius equation shows, in agreement with our results, that in such cases where there is a rise in temperature of 10K, the rate constant and hence the rate of reaction will approximately double. Thus a reaction taking place at 323K should therefore go 23 or8 times faster than the same reaction at 293K. This can be illustrated in the distribution of the kinetic energies of gas molecules at T K and (T+10)K, whereby the amount of molecules with sufficient energy to react at (T+10)K is double that at T K-illustrated by the shaded area, thus confirming my results:
With regards to the effect of surface area, my results, in accordance with my predictions have shown that an increased surface area decreases the duration of the reaction and hence increases the rate of reaction. The graph showing rate of reaction against surface area shows a positive correlation. The accuracy of the results may be questioned by the estimation of the surface area of the powdered calcium carbonate. From the graph of surface area against time we may approximate it to be 23cm2. However this seems somewhat unrealistic as it is likely that it has a much larger surface area. This error may be explained by the series of experimental errors, in particular the fact that the initial surface area of the four pieces of calcium carbonate were approximated to be the same. A more reliable set of results may have been obtained had the pieces all had exactly the same shape and thus surface area.
By comparing the different results we find that changing the calcium carbonate from solid to powder form had the greatest effect on the rate of reaction, it increased it by a factor of 22. To match this we would have had to increase the temperature by 40K or increased the concentration by a factor of 22. Changing surface area was probably the easiest and most practical as well as the most effective.
On a whole we may say that our experiment has proved that increasing surface area, concentration and temperature all increase the rate of reaction. However, we may only say this regarding the reaction that we have considered. Since, although we may have found that increasing these factors increases the rate of reaction this may not necessarily be true for all reactions. Indeed, as we mentioned, in reactions of first order with respect to one of the reactants, altering concentration has no effect. Also in enzyme catalysed reaction, above a certain temperature, increase in temperature in fact has a negative effect on the rate of reaction.
Provided that the equipment needed is available, we may repeat the experiment using different methods of analysis. The conductmetric or measuring gas evolved methods, as explained, may be worth trying. Also, different types of reactions may be considered. For example in considering the effect of temperature, we may consider explosive reactions or enzyme catalysed reactions, a graph of reaction rate against temperature would look something like as follows:
most reactions explosive reactions enzyme catalysed reactions
It may also be of some use to consider other factors that effect the rate of reaction such as catalysis and light. A catalyst is a substance which alters the rate of reaction without itself undergoing any permanent change. This is as a provides a new reaction path for the breaking and rearrangement of bonds, a path that in most cases has a lower activation energy. Using a basic analogy, we may say that a catalysed reaction is similar to a pole-vaulting event in which the bar has been lowered and hence many more athletes get over it. Light too has an effect on certain reactions. Photosynthesis and photography are two common examples.
Bibliography
Atherton, M., An Experimental Introduction to Reaction Kinetics
Benson, S.W., The Foundations of Chemical Kinetics
Laidler, K.J., “Just what is a Transition State?”
Laidler, K,.J., Physical Chemistry
Lewis. M., Thinking Chemistry
Liptrot, G.F, Modern Inorganic Chemistry
Murell, J.N. Introduction to The Theory of Molecular Collisions
Ramsden, E.N., Chemistry for G.C.S.E
Urry, D.W., Henry Errying
Detonation is the word that first referred to the noise of an explosion, but which in the latter part of this century has been used to describe the explosion itself. The process of a detonation results in a shock wave being formed, at which there is an extremely rapid rise in pressure and temperature.
Rusting refers to the oxidation process that takes place with certain metals. An example being Iron. A flaky, reddish brown substance best regarded as a hydrated iron(III) oxide is produced. This product is undesirable and costs Britain an estimated £2 000 000 000 per annum.
Haber Process: Industrial production of ammonia -reversible reaction. High pressure, temperature and the presence of a substance known as a catalyst which has the effect of reducing the energy pathway or activation energy needed to start the reaction, Iron in this case, are ways in which the rate of this reaction is increased.
Archaeological dating is the process of dating the remains of living things from the amount of carbon-14 in them, based on the simultaneous production and disintegration of this radioactive isotope of carbon.
An ion is a charged particle
Svante August Arrhenius was a Swedish chemist who lived from 1859-1927 and is one of those who helped found modern chemistry. His awards and honors include the 1903 Nobel prize in chemistry. His theory of electrolytic dissociation eventually became the cornerstone of modern physical chemistry.
A nanometer is 10-9 metres.
Henry Errying (1901-1981) was born in Colonia Juarez, Mexico, of American parents. In 1927 he obtained a Ph.D. degree in physical chemistry. Much of Errying’s research was in chemical kinetics but he also worked on the kinetics of physical processes such as diffusion and on the structures of solids and liquids. Perhaps his most important contribution was his formulation, in 1935, of the transition-state theory. Yet his failure to win a Nobel prize has been a matter of surprise among many physical chemists.
The order of reaction as stated above should not be confused with the overall order. References to order of a reaction are regarding what is known as partial order, or the order of reaction with respect to a particular reactant. The overall order is the sum of all the partial orders.
Brownian motion refers to the constant random movement of tiny particles in a fluid or gas. The phenomenon was first discovered in 1827 by the British botanist Robert Brown. Brown also made other notable scientific contributions including the discovery of the nucleus of the vegetable as well as the discovery of the distinction between angiosperms and gymnosperms.
The absolute or Kelvin scale which was invented by the British mathematician and physicist William Thomson, and is most commonly used in scientific work. In this scale ,zero K corresponds with absolute zero or -273.16° C, with degree intervals similar to those on the centigrade scale.
Jacobus Hendricus Van’t Hoff (1852-1911) was a Dutch physical chemist and Nobel laureate. Born in Rotterdam he became professor of chemistry, mineralogy, and geology at the University of Amsterdam in 1878. His theory that was put forward to explain the structure of organic chemistry has won him the title by some as the father of physical chemistry. In 1901 he was awarded a Nobel Prize in chemistry for his work relating thermodynamics to chemical reactions.
