Safety
A pair of safety goggles will be worn during the experiment to prevent any of the chemicals being spilled into the eyes. An apron will also be worn to prevent my skin or clothes from being harmed.
Fair Test
In order for my results to be valid, the experiment must be a fair one. I will use the same procedures each time for judging when the ‘X’ disappears. I will make sure that the measuring cylinders for the HCL and Na2S2O3 will not be mixed up. The amount of HCL will be the 100ml each time, same for the Sodium Thiosulphate; 100ml which will be constant. All of these precautions will make my final result more reliable and should keep anomalies to a low minimum. Thus making the final conclusion accurate and reliable.
Prediction
I predict that as the concentration of Sodium Thiosulphate is increases, the rate of reaction will increase. This means that my graph should have positive correlation, and will probably be curved, as the increase in rate of reaction will not be the same as the concentration increases. (See fig.1). All this can be justified by relating to the ‘collision theory.’ As the concentration of a chemical is increased per unit volume, the more chance there is of a collision with greater energy. Particles with more energy are likely to overcome the activation energy barrier to react, thus reacting successfully. All this can be understood better with a full understanding of the ‘collision theory’. For a reaction to occur, particles have to collide with each other. Only a small percent result in a reaction. This is due to the energy barrier to overcome. Only particles with enough energy to overcome the barrier are called the ‘activation-energy,’ or Ea. The size of this ‘activation-energy’ is different for different reactions. If the frequency of collisions is increased, the rate of reaction will increase. However, the percent of successful collisions remains the same. An increase in the frequency of collisions can be achieved by increasing the concentration, pressure, or surface area. (I will discuss the concentration, as this is more relevant, while the others aren’t).
Secondary Sources
A reaction can be made to go faster or slower by changing the concentration of a reactant. Everything is kept the same each time, except the concentration of the acid.
A reaction goes faster when the concentration of a reactant is increased. For the reaction below, the rate doubles when the concentration of acid is doubled.
In order for a reaction to happen: -
▪ they must collide with each other
▪ the collisions must have enough energy
If there are lots of successful collisions in a given minute, then a lot of Sodium is produced in that minute. In other words, the reaction goes quickly – its rate is high. If there are not enough, it’s rate is low.
The rate of reaction depends on how many successful collisions there are in a given unit of time.
“Co-ordinated Science.
Chemistry: Second Edition
Gallagher, Ingram & Whitehead”.
Reaction Rates
Chemical reactions require varying lengths of time for completion, depending upon the characteristics of the reactants and products and the conditions under which the reaction is taking place. Chemical Kinetics is the study of reaction rates, how reaction rates change under varying conditions and by which mechanism the reaction proceeds.
What factors affect the rate of reaction?
We’ve already said that the characteristics of the reactants affect the rate of the reaction, what I want to do here is see what physical factors affect the rate. I can list these;
1: The concentration of the reactants. The more concentrated the faster the rate (note in some cases the rate may be unaffected by the concentration of a particular reactant provided it is present at a minimum concentration). Remember for gasses, increasing the pressure simply increases the concentration so that’s the same thing. This can be named as an Independent/Dependant variable, as I can change the concentration. The HCL will be a Dependant variable while the Sodium Thiosulphate will be an Independent variable).
2: Temperature. Usually reactions speed up with increasing temperature (“100C rise double rate”). This can be named as a Dependant variable, as I have no control what so ever on the temperature. However, I will not be testing the temperature.
3: Physical state of reactants. Powders react faster than blocks – greater surface area and since the reaction occurs at the surface we get a faster rate. This factor does not apply to my investigation.
4: The presence (and concentration/physical form) of a catalyst (or inhibitor). A catalyst speeds up a reaction, an inhibitor slows it down. This factor can be named an Independent variable as I can add a catalyst, but I will not be doing so.
5: Light. Light of a particular wavelength may also speed up a reaction. E.g. alkanes reacting with Halogens. Light can be names a Dependant variable, as I have no control over the wavelength of light.
I’ve used the term reaction rates, so I should define what we mean by this; the reaction rate is the increase in molar concentration of product of a reaction per unit time or the decrease in molar concentration of reactant in unit time.
Rate Law and Reacting Order
If we examine the effect on the rate of a reaction by changing the initial concentration of reactants, we may be able to derive the rate law and hence the reaction order.
Consider the following reaction: -
2NO2 + F2 → 2NO2F
If the concentration of NO2 is doubled then the rate is doubled, likewise when the concentration of the fluorine is doubled the rate doubles and so we get the following rate law;
Rate = k [NO2][F2]
The rate law is an equation that relates the rate of a reaction to the concentration of reactants raised to various powers.
The rate constant, k, is proportionally constant in the relationship between rate and concentrations. This has a fixed value at any given temperature but varies with temperature.
First order reactions
Consider a first order reaction: -
A → Products
The rate law is of the form: -
Rate = k [A]/t
If we set a = the initial concentration of A, x = loss of A with time, then (a-x) is the concentration of A at any given time.
‘www.newi.ac.uk’
Fig.1
Using the results shown in ‘fig.2’ I am going to construct a graph according to the information. I will draw the line-of-best-fit, along with notes on any obvious trends.
Analysis
Referring to fig.3, the main points to notes about the curve are: -
1: There are no particles with zero energy.
2: The curve doesn’t touch the x-axis, because there were no particles with zero energy – proving that each particle will always have energy.
3: The peak of the curve indicates the most possible energy.
The activation energy for a given reaction can be marked on the distribution curve. Only particles with enough energy equal or greater than the activation energy can react when a collision occurs.
Conclusion
Using the above results, it seems that my prediction was correct. As when the percentage of Sodium Thiosulphate was increased, the time taken decreased. The time more or less doubled (decreased) each time the percentage of Sodium Thiosulphate was increased. However, my results were not entirely accurate – most of the results recorded didn’t sit on the line-of-best-fit.
However, these results alone cannot provide firm proof, which could draw an accurate conclusion. Looking back at my prediction, according to what I said; the results should ‘keep-on’ increasing – however, the results do not fit my prediction. Therefore, my results are not firm enough to support my prediction.
Therefore, I will do a pilot study in which I will make a few adjustments that will provide a firm evaluation.
Improvements
Looking back at Fig. 3, the graphs shows that at least one anomaly occurred in the experiment. Fig. 3 shows that there was only one anomaly, which did not fit the pattern – the result increased more than the pattern did. However, each of the results also didn’t fit the line-of-best-fit.
The reason being that the use of the equipment must have been mistreated, or the procedure was not followed accurately either at the start or between the experiments.
I can’t say I didn’t expect this occurrence even know I tried to avoid it. When I was judging whether the ‘X’ had disappeared sufficiently, time was going on and I must have been complacent when judging the ‘disappearance’.
Another investigation should be conducted but with greater care and planning, such as having a light intensity meter and a light bulb to judge accurately when the ‘X’ disappeared properly, as my results would have been the same fairness throughout. Hence, doing this would definitely cancel out any of these anomalies and turn them into proper results. I could also use larger concentrations, which should decrease the time taken further thus providing more solid results, which would help my conclusion.
I am now going to do a Pilot study in which I will learn from my errors previously and redo the investigation. I will use the same apparatus, method and knowledge, (I will also not include a prediction due to the wildness nature of the results) - to keep my entire investigation fair. Changes that will be made in order to provide a more solid set of results and conclusion are: -
1: Increasing the concentration of Sodium Thiosulphate to 3.0 - 4.5. Using moles instead. To see if this will increase the rate of reaction.
2: Use of HCL at 2 moles instead of 0.5 moles.
3: I will also calculate the time taken for each mole to react.
4: Calculate the rate of reaction as well, 1/t.
From these variations I will be able to collect more accurate results along with a firmer evaluation than my preliminary study.
Pilot Study
I have already stated the changes that will be made in order to gather more detailed information for my evaluation.
My results
Time Taken / Mole
I am going to compare the lines-of-best-fit on the graph. This shall give me more information for my evaluation.
Fig. 6 = 136 (y) ∕ 4 (x) = 34m/s
This results shows that for every second during the reaction, 34 moles reacted. I thought that the reaction would react more moles than this result.
Evaluation
In this experiment I have that as the concentration is increased, the time taken for the reaction to take place decreases. This means the rate of a reaction increases as it takes less time for a reaction to take place, so more take place per second as shown by the above result. However looking at fig. 6, it shows that when the amount of Sodium Thiosulphate (M) decreases the time taken increases.
Using fig. 6 and 7, with line-of-best-fit, I can draw a more solid conclusion than that of my first one. Firstly, fig. 6 shows positive correlation, meaning as the concentration increased, the time taken for a reaction to take place decreased. Still looking at fig. 6, it shows that the concentration against time taken was positive correlation. As when the concentration decreased, the time taken increased. This is because when the concentration was increased, there are more particles per unit volume – resulting to more collisions. Particles with more energy are likely to overcome the activation energy barrier to reaction, and thus reacting successfully. The total amount of moles reacting per second is shown above.
Fig. 7 shows that when the concentrations were relatively low (0.5M – 1.0M), the increase of rate x 1000 was also fairly small (increasing from 5.1 to 6.7). There was then a gradual increase in the difference, and between 1.0M and 1.5M the rate more than doubled from 6.7 to 10.2 compared to previous results. This shows that there are far enough collisions at 4.5M than at lower concentrations.
For this to make sense, it is necessary if I recap on the collision theory: For a reaction to take place, particles must collide with each other, only a small percent result in a reaction. This is all due to the energy barrier to overcome. Any particle with enough energy to overcome the barrier is called the activation energy, or the Ea. The size of this activation energy is different for different reactions. If the frequency of collisions is increased, the rate of reaction will increase. However, the percent of successful collisions remains the same. An increase in the frequency of collisions can be achieved by increasing the concentration, pressure, or surface area.
Like stated previously, most of the results that had been recorded didn’t fit the line-of-best-fit. Which could be a result of inaccurate calculation, temperature or even the concentration levels. Looking at the curve on fig.3, it shows the intake of heat energy that results in a faster reaction time – this supports the fact that this reaction is an endothermic reaction and when heat energy is absorbed, the time taken decreases substantially. However, it seems that the HCL had reached its factorial limit, which would have affected the time recorded – therefore, higher concentrations should have been used. This is proved when the results slope horizontally on fig. 3 & 6.
