Successful collision frequency α concentration
Rate of the reaction α concentration
Surface area of the reactants
The surface area of the reactants affects the rate of a reaction, also because of the collision theory. The bigger the surface area of the reactants, the smaller each piece will be, and the more pieces of reacting material there will be. With more pieces, more reaction sites are created, and so there will be more sites where collisions can take place, increasing the collision frequency. With an increase in the collision frequency, an increase in the successful collision frequency results and so the faster the reaction accelerates.
The diagram above shows that when the reactants are big pieces (left) there will only be a small surface area where a reaction can occur, resulting in less particles being able to collide and therefore a smaller collision frequency. When the reactant is crushed into many smaller pieces (right) however, the surface area is greatly increased. Thus, resulting in more particles being able to collide and therefore a much higher collision frequency.
This idea is made clearer in the diagram above, where with a lower surface area (left), there is room for 8 particles of reactant 2 to collide with reactant 1. With the same amount of reactant 1, however broken into smaller pieces, creating a larger surface area (right), room for 16 particles of reactant 2 is made.
More room is made for collisions; therefore more collisions will be made. Thus, a greater successful collision frequency will be made, increasing the rate of the reaction.
The collision theory states that with an increase in the collision frequency, the greater the successful collision frequency will be and so the greater number of products forming, increasing the rate of the reaction.
We can therefore state that the surface area of the reactants is proportional to the collision frequency:
Surface area α collision frequency
The successful collision frequency is also proportional to the collision frequency:
Successful collision frequency α collision frequency
Therefore, we can state that the surface area is directly proportional to the successful collision frequency and so the rate of the reaction:
Successful collision frequency α surface area
Rate of the reaction α surface area
Temperature as a factor
Hypothesis
I predict that an increase in temperature will cause an increase in the rate of the reaction.
I base my prediction on the collision theory. It says that for particles to react with each other, they must collide. Even if they collide, however, they must meet a certain energy level in order for it to be a successful collision and so a reaction taking place. If they do not meet this energy requirement (activation energy), then the collision will be unsuccessful and no reaction will occur.
With an increase in temperature, the average kinetic energy of the particles is increased and so there will be a greater chance that the particles colliding will have met the activation energy requirement.
I also base my prediction on the kinetic theory, where particles are excited more as they gain kinetic energy. This causes them to move faster and by doing so collide more often. With more total collisions, naturally comes a greater successful collision frequency, and so more reactions occur, increasing the rate of the reaction.
I further predict, using the secondary source below, that an increase of 10°C will double the rate of the reaction.
“The effect of temperature on rates of reaction is important. Ordinarily, raising the temperature 10º (Celsius) approximately doubles the rate of most but not all reactions. This is because the chance that molecules are sufficiently energy rich to have achieved activation level and undergo reaction is greatly increased by an increase in temperature.”
Source: Britannica Online
Surface area of the reactants as a factor
Hypothesis
I predict that as the surface area is increased, the rate of the reaction will increase also.
My reason being that during a chemical reaction, molecules of the liquid reactant collide with molecules of the solid reactant. However, the liquid molecules can only collide with the solid molecules at the liquid-solid interface, i.e. only the surface of the solid.
By increasing the surface area of the solid, there will be more room for collisions to take place between the liquid molecules and the solid molecules, at the liquid-solid interface. Thus, there will be more collisions, and therefore a greater successful collision frequency also, increasing the rate of the reaction.
Preliminary Experiment
I have also done a test between magnesium ribbon and hydrochloric acid. I reacted 1cm of Mg strip with 20cm³ of the acid. I then repeated the experiment, halving a new strip of 1cm of Mg each time, to create larger surface areas. These are my results:
As you can see from the table, as the surface area increases, the rate of the reaction does also. I can use this information to back my prediction, that the bigger the surface area, the faster the rate of the reaction. I can also use it when processing the results obtained from this investigation, to aid me in deducing a firm conclusion
Concentration of the reactants as a factor
Hypothesis
I predict that as the concentration of the reactants is increased, so will the rate of the reaction linearly increase.
I base my prediction on the collision theory. As the concentration of the products is increased, there is more of the product colliding, and so more collisions per second. With an increase in the total collision frequency comes an increase in the successful collision frequency and therefore more reactions happening per second. Thus, the rate of the reaction is increased.
I also base my prediction on the kinetic theory. Throughout the particles of the reactants, there are various energy levels. In order for a collision to be successful, the energy level of the colliding particle must meet the activation energy required for the reaction to take place.
By increasing the concentration of the reactants, you increase the total collision frequency, thus increasing the frequency of those collisions that exceed the activation energy requirement, which lead to successful reactions taking place.
In addition to the above, I can also predict the volume of Hydrogen gas that should be made in the experiment. Because each of the different concentration reactions contain the same volume of acid and same weight of Magnesium, the amount of Hydrogen gas produced will be uniform. Therefore, when increasing the concentration the speed of the reaction will increase and not the amount of products produced.
I can use quantitative chemistry to help me predict the amount of hydrogen gas that will be produced in the reaction. This will be done via finding out how many “moles” of reactants there are, and therefore how many “moles” of products that be should be obtained in theory. Consequently, when finding out how many “moles” of Hydrogen gas is produced, another formula can be used to find out the exact volume of Hydrogen that will be produced.
A is simply a unit of measure that is usually used to measure very tiny objects like atoms and molecules. In comparison, one dozen is a measure in quantities of 12. But, one mole is a measure in a quantity that is much larger. Again, a dozen is 12 objects but a mole is 602,200,000,000,000,000,000,000 objects. This is a huge number and easier and quicker to write using scientific notation or 6.022x1023. This number is also called number. One mole of atoms of every element is equal to the atomic mass of the element.
I can use this notion to work out how many moles of Magnesium (Mg) I will be using in the reactions.
The symbol equation for the reaction is as follows:
Mg (s) + 2HCl (aq) → MgCl2 (aq) + H2 (g)
1 Mole 2 Moles 1 Mole 1 Mole
0.046g 25cm3 46.1cm3
The formula for finding out the number of moles in a given mass is as follows:
Mass of Magnesium = 0.046g
Mr (Relative atomic mass) of Magnesium = 24
Therefore,
Moles= 0.046/24=0.00192 (3 S.F.)
The equation above shows that 1 mole of Magnesium reacts with 2 moles of Hydrochloric acid to produce 1 mole of Magnesium Chloride and 1 mole of Hydrogen molecules.
The ratio between Magnesium and Hydrogen is 1:1. Therefore, the number of moles of Magnesium to be reacted will be equivalent to the number of moles of Hydrogen molecules produced.
From the working above, it has become evident that I will be using 0.00192 moles of Magnesium in my reactions. Therefore, an equal number of moles of Hydrogen gas will also be produced; therefore, 0.00192 moles of Hydrogen gas will be produced in the reaction.
“For reactions involving gases reacting (and any gaseous products that form), it was observed that they do so in volumes which are in ratios of small whole numbers, when measured under the same conditions of temperature and pressure. This is expressed as Gay-Lussac's Law (1808). This observation was explained by Avogadro, and is expressed as Avogadro's Law (1811). Equal volumes of any gases measured under the same conditions of temperature and pressure contain equal numbers of molecules (or atoms if the gas in monatomic).”
Secondary Source: Britannica Online
The Molar Volume is the volume of 1 mole of a gas. This, of course, depends upon the temperature and pressure at which it is measured. The Molar Volume at room temperature and pressure (298 K, 1 atm.) is 24 dm3.
From this, I can work out, in theory, how much Hydrogen gas will be produced from the reaction.
The formula for finding out the volume occupied by a gas at room temperature and pressure is as follows:
Moles of Hydrogen gas produced= 0.00192 (3 S.F.)
Volume of gas taken by 1 mole of gas atoms or molecules = 24 dm3
Therefore,
Volume= 0.00192 * 24000 cm3 = 46.1cm3 (3 S.F.)
I know, therefore, that I will be expecting to obtain a volume of roughly around 46cm3 of Hydrogen gas (H2) from each of the different concentration reactions.
Preliminary Experiment
I tested my theory at home.
I got two “Rennie” tablets and placed one in a cup of lemon juice, and the other in a cup of diluted lemon juice, thus creating an independent concentration variable. The tablet reacted faster in the cup of concentrated lemon juice, rather than the diluted one.
Therefore, I could conclude that the concentration of the reactants does increase the rate that the reaction commences.
I can use these results to back my prediction, and also to aid me in my analysis of the results obtained from this investigation.
In addition to doing this, I also carried out a further preliminary experiment using Magnesium and six different concentrations that I will investigate later. I didn’t record any measurement at this stage, because I was only analysing how fast the reaction rates were, by looking at the amount of water that was being displaced by the hydrogen gas migrating into the tube. This is called collecting gases “over water”
It is the best method of collecting a gas so long as the gas is insoluble or slightly soluble in water. This method can be used for insoluble gases which are more dense or less dense than air e.g. hydrogen, nitrogen, methane and oxygen.
This allowed me to make important decisions about the time intervals that will be of use to me in my official experiment for recording the amount of Hydrogen gas being evolved.
Variables
Independent
- Concentration of reactant (Hydrochloric Acid);
Control
- Length (and therefore mass) of Magnesium Ribbon;
- Volume of Hydrochloric acid used (25 cm³);
- Surface area of conical flask (this must be kept the same)
Apparatus
- Safety goggles
- Conical flask
- Gas Syringe
- Stopwatch
- Measuring cylinder
- Hydrochloric acid
- Magnesium Ribbon
- Water
Method
- Range of measurements: 2M, 1.75M, 1.5M, 1.25M, 1.0M HCl acid
- Measurements repeated: thrice times necessary. In the event of varied results, experiment is repeated as appropriate until sustainable results with little variations is obtained;
- Volume of the acid used: 25 cm³
For each of the concentrations, the following method was used:
After measuring out the appropriate concentration of acid (which is a control volume of 25 cm³) with a measuring cylinder, I poured it into the conical flask. I made sure that the gas syringe that I was using was set to zero (therefore, no gas inside to begin with). Then, I inserted one strip of Magnesium into the conical flask (where the acid is) and immediately placed the bung on top, to ensure that no Hydrogen gas escapes. By this time of course, my colleague ensured that he started the stop watch as soon as I had placed the Magnesium inside. I took measurements from the gas syringe every certain interval of time.
My preliminary work helped me to choose the time intervals for measure, as I had examined how much hydrogen gas was being produced from the upward delivery method. The 1.5M and 1.0M, produced relatively small amounts of hydrogen gas as the volume of water in the tube was moving down very slowly indeed. I deduced that it would be more appropriate to reduce the intervals from 2 seconds to every 10 and 30 seconds respectively.
I repeated each experiment a further two times, and took their recordings. I then took an average result for increased accuracy. If one of the results achieved from the two repeat experiments varies greatly, the experiment shall be repeated for a fourth time to hopefully give three sets of reliable results. Or obviously, if both of the repeat experiments vary greatly from the original experiment, then the experiment will be repeated a further two times and as necessary, to ensure that three sets of close and reliable results are obtained.
Safety Precautions:
As throughout these experiments, the use of acid was involved, I wore protective eye safety goggles as a safety precaution. Also, I carried out the tests along with a colleague at the back of the lab, away from any obstacles. I put all other unnecessary stuff away, which can pose as a potential hazard by, for example, obstructing pathways. This included tucking stools into their appropriate places and putting things like my coat and school bag away. I also made sure that I used the apparatus as safely as possible to avoid any accidents.
Results
The investigations I carried out were successful. The results were suitable and generally quite good. A fellow colleague aided me in carrying out the experiment and making recordings of measurements.
To ensure that I obtain the most accurate results as possible, I carried out each of the different concentrations 3 times. If one of these results proved to be very different to the others, the experiment will be repeated for a fourth time. If, again the result seems to be different, the experiment would be repeated up until 3 sets of similar results are obtained.
Experiments that produced results which have been discarded are columns colour coded GREEN. The end times for each of the experiments are GREY, and the average of the three correct trials have been recorded in the Average results column colour ceded DARK ORANGE.
I took care with the safety precautions and carried out the investigation safely.
I am now displaying the results I obtained for each of the concentrations in tabular form:
2M Concentration (HCl)
1.75M Concentration (HCl)
1.5M Concentration (HCl)
1.25M Concentration (HCl)
1.0M Concentration (HCl)
I drew up these 5 concentration results in graphical form (volume of gas produce against time taken) so that I could use them when deducing a conclusion. Furthermore, I drew up the average of the time taken for each concentration (as summarised in the table below) as well as an inverse of time for each of the concentrations. As well as that, I drew up another graph to illustrate the initial rate of reaction of each of the concentrations.
Analysis: Concentration of the reactants as a factor
The aim of this investigation was to find out how changes in the concentration of the reactants change the rate at which the reaction accelerates. I completed my experiments successfully, obtaining results that I can now use to deduce a firm conclusion as to how variations in the concentration of the reactants affect the rate of the reaction. My results show that as the concentration of the reactants is increased, the time taken for the reaction to be completed decreases.
I have drawn up a line-of-best-fit graph to graphically display my results. From the graph I can notice a trend; the higher the concentration, the less time taken for the reaction to complete.
The graph also shows that at the lower concentrations, there is a sharp decrease in the time taken for the reaction to complete, whereas this sharp fall tends to level off towards the higher concentrations.
The reason for this fall in the time taken for the reaction to complete is due to the collision theory. This states that when the concentration of the reactants is increased, the more particles there will be, and thus, the grater the collision frequency will be. With an increase in the collision frequency naturally comes an increase in the successful collision frequency as well, and so more reactions will take place per second. Thus, it will take less time for the whole of the reaction to complete. Therefore, an increase in the rate of the reaction is made.
The graph however begins to level off towards the higher concentrations. Although due to the collision theory, further increases in the concentration will increase the rate of the reaction, the rate at which the reaction accelerates at this point is limited by other factors such as the surface area at which collisions can take place. Although there are more particles ready to collide and react at these higher concentrations, there may be not enough room on the surface area of the second reactant for them to collide with. Therefore, these extra particles ‘wait’ around the active reaction site, ready to collide as soon as space is made available for them. Thus, the reaction rate still increases, but not as fast as it would have done at the lower concentrations.
Therefore, the results obtained support my prediction that as the concentration of the reactants is increased, so will the rate of the reaction linearly increase.
The results obtained from my preliminary test also confirm the accuracy of the results obtained in this investigation, as well as support my prediction.
The table above gives the volume of Hydrogen gas collected from each of the different concentration of HCL acid. It can be seen that the average result of 45.7cm³ Hydrogen gas is very close to the predicted volume of gas via using quantitative chemistry (46cm³), and therefore goes to show that the experiments were done very accurately indeed.
Evaluation
The investigation went quite successfully, with some very good results obtained. These results I found were quite sufficient to base a firm conclusion on. The results were sequenced, with a visible trend displayed on the graphs. There was a little percentage error, and no substantial anomalous results were obtained.
The repeated measurements increased the overall accuracy of the results and the general procedure used was quite good, enabling the output of quite accurate results. Nonetheless, various steps could be made to the procedure to ensure the results obtained would be even more accurate.
I suggest using a computerised data logging set of apparatus where the exact starting masses can be recorded by the computer. Also, graphs of the reaction rate would be automatically created, and these could aid me further in my conclusions.
It would also be interesting to investigate how other factors affect the rate at which the reaction occurs. Such factors could include the effect of catalysts on the rate of the reaction, as well as the effect of the pH on the rate at which a reaction occurs. These are in addition to all the other factors I listed in my introduction, in which I would like to test them out to see if they correspond to the theory and the preliminary experiments I did for each of them.
