Rates of Reaction
Group Members: Alishan Artani, Joshua Li-Taylor, Lageishon Mohanadas, Zachery Oman
Given Date: October, 11, 2008
Due Date: November, 30, 2008

I. Introduction

Throughout the study of chemistry, chemicals have been discovered as well as created as a result. For example, tin and copper are substances that have been found, in solid state, in the ground. But when combined, produces bronze, a metal with stronger bonds. This is called a chemical reaction. Chemicals reactions usually occur over a certain duration of time when mixing chemicals. When comparing the speed at which two or more chemicals combine to create a product, we use reactions rates. Reactions rates are true as long as a change occurs and as long as the value of a reaction rate is positive, because a negative value would imply that change has occurred before mixing the chemicals. Asides from these two situations a reaction rate can vary from any duration of time. Over time, information about reaction rates has been gathered and four variables, that can have an influence on reactions rates, have been considered. They are the difference in surface area, the difference in concentration, the difference in temperature and if a catalyst is added. These variables are explained in further detail below and are supported by the collision theory proposed by Max Trautz. To test if these four variables affect the reaction rate in chemical reactions, we will perform an experiment for each variable. Our hypothesis is that all these factors are true and all of them must been considered when comparing reactions rates, especially when comparing the results of the same experiment. Under these situations precision is needed and forgetting these variables can cause a huge margin error depending on the requirements of the experiment.
Surface Area, one of the four variables that affect reaction rates, considers the effect a solvent has to a solute in when there are different quantities but has the same mass. It is believed that that the solute with many but smaller pieces will dissolve faster then one larger piece. For example, a sponge dropped on a puddle will absorb the water in a reasonable short time, but several pieces of that same sponge will absorb the puddle faster. This is because the solute is exposed at more sides when in pieces, providing more opportunities for collision rather than a huge piece. This is based on the collision theory that states that in order to achieve a successful reaction, the reactant particles must collide with each other in a specific angle and have the certain energy level. Activation energy is the minimum amount of energy required for these collisions. The ones that collide in this certain manner and have the energy create the product by destroying their bonds and creating new bonds. If the chemical were to be exposed to more sides and have more pieces there would be a higher chance of the reactants colliding with each other. To demonstrate that this is true, we will be comparing the rate at which CuSO4 dissolves into a 50mL beaker of water in three different sizes with the same mass.
Concentration is the second factor when observing reaction rates because it deals with the difference of particles in the same named chemical with a common mass between the two. A reactant that can mix to form a product will do so in a certain amount of time, but the same reactant at a higher concentration will mix and form the same product but faster. According to the collision theory this is because there is a higher amount of particles in the reactant when it has a higher concentration. Since there are more particles there is a higher chance that the reactant’s particles will collide with each other. We will see if this is correct by using 2g of crystallized Na2S2O3×5H2O mixed with 10mL of water and add 10mL of different concentrations of HCl. We will compare results by timing and determining the measuring the difference in times.
Temperature is assumed to be a factor of reaction rates because of the difference in affect to it has to chemicals. For example, having a higher temperature in an equation can produce more heat, which in turn causes the chemical reactions to speed up as it is being heated and the chemical(s) is dissolving faster into the other(s). Having a colder temperature in an equation will have the opposite affect. In terms of the collision theory, reaction rates will speed up because as temperatures increase, so does the amount of energy in the reactants. This energy increases the speed at which the reactant’s particles travel giving the reactants a better chance to have a successful collision. To show this we will repeat the procedures from the concentration experiment but heat the Na2S2O3×5H2O and 10mL of water at maximum and medium heat before adding the HCl. We will compare these results to the results from the concentration experiment where the mixture was not heated.
Catalysts are the last factors that affect reactions rates because they are chemicals that act as a third party and stimulate reactants to create a product without being consumed in the reaction itself. Shown by Boltzmann distribution and the energy profile diagram catalysts work by creating an alternative route for reaction that has a different activation energy and transition state. Transition state is the phase whereby reactants are at the highest level of energy will become products and will not revert back. On this route, catalysts decrease the activation energy. This means is that collisions have a better chance colliding because it will have the energy needed to have a successful collision to reach the transition state. Different types of catalysts can manipulate the equation in different ways like increasing the reaction rate or lowering the overall temperature of an equation. Our experiment to show if a catalyst can produce a reaction involves the combination of potassium permanganate, sulphuric acid and oxalic acid, respectively. Then we compare the results of this mixture when heated and when the catalyst is added.


II. Materials

i) Basic Materials
· Water
· Two Standard beakers
· Two graduated cylinders
· Stopwatch
· Scoopula
· Pipette
· Goggles
ii) Experiment 1: Surface Area
· CuSO4 (copper (II) sulphate) in the following forms:
o Large crystals
o Medium sized crystals
o Powdered
· Mortar and pestle
· Electronic scale
· Magnetic stirrer

iii) Experiment 2: Catalyst
· KMnO4 (potassium permanganate)
· H2SO4 (sulphuric acid)
· C2O4 (oxalic acid)
· Thermometer
· Beaker Tongs
· Hotplate

iv) Experiments 3 & 4: Concentration and Temperature
· Permanent Marker
· White paper
· HCl ≥ 4M (Hydrochloric acid)
· Crystallized Na2S2O3×5H2O (sodium thiosulphate)
· Stirring Rod
· Beaker Tongs
· Hotplate


III. Procedure

i) Experiment 1: Surface Area
1. Using a scoopula, scoop out approximately 2g of large, crystallized CuSO4. Weigh the sample on an electronic scale to verify its mass and then proceed once a correct mass of CuSO4 is obtained.
2. Fill a beaker with 50mL of water.
3. Set up the beaker with a magnetic stirrer on the magnetic plate. Set the speed to 2 and wait for the stirrer to make its rotations consistent.
4. Pour the CuSO4 into the beaker and start the timer immediately.
5. Record the time it takes for the CuSO4 to dissolve completely with no particles left floating in the solution.
6. Repeat steps 1-5 using medium-sized, crystallized CuSO4.
7. Repeat steps 1-5 using powdered crystallized CuSO4 that has been made into a finer powder to distinguish it from the medium-sized crystals. Do this using a mortar and pestle.
8. Clean up work are and prepare equipment for the second portion of the lab.
ii) Experiment 2: Catalyst
9. Mix 3mL KMnO4 with 1mL H2SO4 to create an acidic medium in which the reaction can take place.
10. Mix the product of the previous with 5mL Na2S2O3×5H2O.
11. Record observations.
12. Repeat steps 9-11 under the following two conditions:
a) Heat the first to a minimum of 70°C.
b) Place MnSO4 in the product.
13. Clean up work area and prepare equipment for the third and final portion of the lab.
iii) Experiments 3 & 4: Concentration & Temperature
14. Mark a piece of white paper with an X using a permanent marker and place on the lab bench.
15. Fill a beaker using a scoopula with 2g of crystallized Na2S2O3×5H2O and add 10mL of water. Stir until it dissolves fully.
16. Placing the beaker on the X, pour 10mL HCl 4M into the beaker and begin the timer immediately.
17. Stop the timer once the X is no longer visible when viewing it through the product (originally colourless, clear liquid should turn a greenish-brown, opaque liquid). Note: Be careful not to look directly from above the beaker as fumes are dangerous and should not be inhaled. View the X by looking at it from an angle.
18. Clean out the beaker and equipment and repeat steps 15-17 using concentrations of 10mL HCl of 2M, 1M, and 0.5M.
19. Now, repeat steps 15-17 twice more using only 10mL HCl 1M. This time however, the solution of Na2S2O3×5H2O and water will first be heated at medium heat for 1 minute on a hot plate, then at maximum heat for 1 minute before proceeding to steps 16 and 17.
IV. Observations

i) Experiment 1: Surface Area

Condition of CuSO4 (crystallized)
Time to Dissolve in water (in minutes)
Large crystal chunks
4:56.5
Medium sized crystals
0:37.98
Fine powder
0:04.78


ii) Experiment 2: Catalyst

Condition of mixed reagents
Observation
Under SATP, no temperature change
Brownish-orange fluid
Heated to 70°C
Clear
After MnSO4 added (the catalyst)
Clear

iii) Experiment 3 & 4: Concentration and Temperature

Concentration
Time (in seconds)
4 M
18.36
2 M
22.52
1M
29.45
0.5 M
36.16

Temperature (with 1M)
Time (in seconds)
High heat
4.21
Medium heat
8.58
No heat
29.45


iv) Conclusion
The hypothesis that all four factors increase the rate of reaction is proven to be true. Increase in temperature, concentration, surface area, and catalysts all support the collision theory as the atoms are interacting more. As the temperature is increased the atoms are heated up and therefore have more energy to move around faster and make more contact. In the same way, if the concentration of one reactant in another is increased, the ratio of one reactant to the other will be smaller so that more atoms can interact with the atom of the other reactant. A greater surface area also allows more atoms of one reactant to be exposed to the other, again, supporting the collision theory so that the reaction can occur faster. Lastly, the catalyst reduces the amount of activation energy needed for the reactants to mingle. Thus, the collision theory proves that these four factors increase the rate of reaction.

V. Discussion
Surface area is directly related to rate of reaction because as the particle size decreases, the total surface area of the reactant increases. As a result, the probability of interactions between molecules increases. This was observed to be true in the experimentation, where, as the particle size of CuSO4 was decreased, thus increasing surface area, the time for the reaction to take place decreased meaning that the rate of the reaction had increased. This is a result of collision theory.
However, exceptions to this case exist. Particle size is not the dependant here, it must be made clear that it is surface area. If there was a heap of very fine powder which was reacting with a gas, how would the gas react with the fine particles within the heap? The answer is that it cannot. This means that the surface area of the substance is just the heap’s surface area, one much too large to create a quick reaction. This can be observed with the example of finely ground magnesium powder. A heap of it burns rather slower than a strip of magnesium ribbon.


Concentration is another consideration when considering rate of reaction. There is a rate equation which mathematically verifies it as well. During the lab, as concentration levels of HCl was decreased, the rate of reaction was also decreased. This shows an example of how greater concentrations will produce higher rates of reaction. The rate equation is as follows:
Rate=k[A]a [B]b
In this equation, k is a variable constant which varies from reaction to reaction. This is the actual rate of reaction under optimal conditions (and varies as conditions such as temperature is changed). This number is difficult to solve for and requires multiple trials to produce accurate results. As well, it is not the purpose of our experiment to quantify rates of reaction, rather, just to determine the rates of reaction via qualitative research. As a result of such, it was not included under Calculations. A and B are the concentrations of the two reagents expressed in mol L-1­. Multiple values of such can be added to the equation. The exponents a and b are variables denoting how much A and B, respectively, are proportionate to the rate of reaction. This uses the order of reaction which can only be found through experimentation and not by looking at a chemical equation. If a was 2 and b was 0 (denoting that B does not affect the rate of reaction because it might be water for instance), the rate of reaction would be dependent upon the concentration of the first chemical. If this began as 1, but was then doubled, that would result in a rate of reaction that has quadrupled [e.g. A=2, a=2, Aa=(2)2, rate=k(4)]. In the experiment carried out, A would be the concentration HCl whereas B would be the concentration of Na2S2O3×5H2O. However, as the concentration of B remained constant throughout all trials, it would not affect the change in rate of reaction. However, concentration does not affect rates of reaction when catalysts are at work because they are already working at their maximum capacity, and in certain multi-step reactions.


Temperature is directly related to rate of reaction and is even proven mathematically by the Arrhenius equation, a complex equation that solves for the constant variable k for the rate of reaction equation aforementioned. As the temperature of a substance increases, its kinetic energy also increases. This means that the molecules are now moving faster and are more spaced apart. Temperature supports the collision theory as now, molecules are moving faster, increasing the average energy of the collisions increasing the rate of reaction, and molecules are moving around more meaning that the chance for molecules to collide is also greatly increased. Therefore, temperature is directly related to collision frequency, proving its relation to rate of reaction. Increased temperature also shows that the molecules have increased energy. This increase in energy is also required so that the reaction takes place in the first place. There is an energy level which particles must first obtain to react with another substance. With heat applied, this number of particles with activation energy increases and is shown in the following graph.

This change in rate of reaction was observed during where the heated samples of Na2S2O3×5H2O dissolved in water reacted a lot quicker with HCl than the one at room temperature.


Catalysts are unique instances where a substance can be added to an already occurring reaction to speed up the rate of reaction to its pinnacle of potential energy. This change is independent of all other dependent conditions which affect the rate of reaction. Although in the lab, it was not proven how to determine a catalyst, it was however proven that catalysts do greatly increase the rate of reaction and eliminate the need for changes in other variables that affect the rate of reaction. This is because a catalyst reduces the activation energy required for particles to react to its minimum possible value and was shown in the experiment conducted where heating of the substances to 70°C was required for the reaction to take place whereas adding the catalyst let the chemical reaction take place at SATP with no altered conditions to the reagents.

VI. Sources of Experimental Error

i) Experiment 1: Surface Area
When the experiment is being done with the highest surface area, or the powdered form of copper pentasulphate, many dust particles could have joined the reactant causing inaccuracy in the result. The amount of dust is unlikely to be a source of major error, but it is there nevertheless. This could be prevented if the experiment was to be done in a room where there is not as much movement and if the room is even cleaner.
The sizes of the crystals of copper pentasulphate were not quantified but rather generalized as small, medium, or large. Further calculations could have been done if the experiment was done with cubic crystals with known surface areas.
There was no accurate way to tell whether the copper pentasulphate was completely dissolved or not. It was all done through human speculation. If a device that could detect when a solution was fully formed was used, less human assumptions would have to be made, so results would be more accurate.

ii) Experiment 2: Catalyst
Observations of the final solution are made right after the eyes perceive the reaction while other reactions may be taking place slowly. The solution should be left out for a longer period of time under SATP conditions to see any changes.
Dust particles may reciprocate the effects of the catalyst. As the catalyst would make the reaction occur faster, if dust particles were flying into the experiment, the reaction would minimally slow down. This small disturbance can be avoided if the lab was to be done in a room with minimal movement and an absolutely clean working space.
This portion of the lab could be extended to how catalysts affect reaction rates rather than whether they do or not. An extremely accurate result may allow proving just that, provided that other catalysts and exact proportions are used. This lab only establishes a very flat point but further quantitative research can reflect in-depth reasoning.

iii) Experiment 3 & 4: Concentration and Temperature
When making the solution of CuSO4 and water, tap water was used so the water wasn’t distilled even if it did fulfill the general purpose. The water may have had a higher TDS rating and did not have a pH level of 7. This may have caused minor discrepancies in the results. Distilled water could have been used in order to get absolutely accurate results.
Since tap water was used, the initial temperature of the water was also unknown. The aim was to have the water filled at room temperature but it may have been slightly warmer, causing a higher rate of reaction, or slightly cooler, causing the rate of reaction to be lower. If the water was left out, while covered, it would adjust to the room temperature and the timings would be more accurate.
The stoppage of time when the X beneath the beaker was fully covered was based on completely on speculation. There was no effective way to measure whether the X had been covered so it was determined through human opinion. To prevent inaccuracy, a computerized sensor that detected the black X could be used, and when it could not detect it anymore, it would indicate so.
When making the solution of CuSO4 and water, tap water was used so the water wasn’t distilled even if it did fulfill the general purpose. The water may have had a higher TDS rating and did not have a pH level of 7. This may have caused minor discrepancies in the results. Distilled water could have been used in order to get absolutely accurate results. The stoppage of time when the X beneath the beaker was fully covered was based on completely on speculation. There was no effective way to measure whether the X had been covered so it was determined through human opinion. To prevent inaccuracy, a computerized sensor that detected the black X could be used, and when it could not detect it anymore, it would indicate so.
During all experimentation, no irregular conditions appeared to be present therefore, the assumption was made that all labs were done under SATP conditions. However, if a lab was completed under abnormal conditions, rate of reactions might have differed, especially if there was a production of a gas from any of the reactions and pressure varied. We could correct this by verifying SATP by measuring the temperature and barometric pressure in the room before proceeding with any experiment.