Measuring Reaction Rate Using Volume of Gas Produced

Chemistry 12 Formal Report Measuring Reaction Rate using Volume of Gas Produced By John Doe 23th October 2012 Introduction An essential element of chemistry is finding reaction rates. This is because chemists need to know how long a reaction should take. In addition to needing to know the rate of a reaction at any point in time to monitor how the reaction is proceeding. Many factors effect reaction rates, two shown above include temperature and concentration. Concentration affects the rate of reactions because the more concentrated a solution the more likely collisions between particles will be.
This is simply because there are more particles present to collide with each other. When the temperature is higher, particles will have more energy. This means that more reactions will happen for two reasons, firstly more particles will come into contact with each other because they are moving around more and secondly because the reactions occur at higher speed making it more likely to succeed. A few other factors are the surface area and if a catalyst is present. The larger the surface area the more collisions will occur because there are more places for molecules to react with each other.
A catalyst affects the rate of reaction not by increasing the number of collisions, but by making more of the collisions that do occur successful. Ordinary household bleach is an aqueous solution of sodium hypochlorite, NaClO, this contains little more than 5% NaClO by mass. Bleaching is caused by the ion. Under normal circumstances this ion breaks down slowly giving off oxygen gas and the chloride ion, . In order to speed up this reaction a catalyst is needed. In this experiment the catalyst used was cobalt (II) nitrate solution.

When this is added to the bleach a black precipitate of cobalt (III) nitrate is formed which acts as a catalyst for the decomposition of The purpose of this experiment was to determine how concentration of reactants and temperature affect the rate of the reaction between bleach and 0. 01M cobalt (II) nitrate solution. In this experiment the volume of gas produced shows the rate of the reaction. Procedure Figure 1 Firstly, all safety protocols were ensured and applied (lab apron and safety goggles). The apparatus was set up with reference to figure 1 above.
Then, the eudiometer was filled with water and inverted into the trough, which was half filled with water. It was held in a vertical position with the burette clamp attached to the stand. The rubber tubing was joined to the top of the glass tube, which goes through the stopper on the flask. The other end of the tubing was then placed into the neck of the eudiometer. 15mL of bleach solution was measured into the 25mL-graduated cylinder and poured into the Erlenmeyer flask. As followed, 5mL of 0. 10M of cobalt (II) nitrate solution was measured and poured into the 10mL-graduated cylinder.
Once ready, the cobalt nitrate solution was poured into the flask containing the bleach solution, and the rubber stopper was immediately slotted in. It was then mixed and stirred as well as recorded (time). It was noted that a black precipitate of cobalt (III) oxide was forming, and from then on the flask was stirred gently and constantly. This was significant to dislodge bubbles of oxygen from the surface of the Co2O3 catalyst. Another thing that was important to note was that if the swirling was stopped or reduced, the rate decreases, so therefore the amount of swirling must be kept steady and uniform throughout the runs.
The total volume of oxygen that had been collected was recorded every 30 seconds until a volume of 50mL was obtained. Also, the actual elapsed time of when the 50mL mark was reached was recorded. Once the first run was successful, the following needed to be repeated the same way: the same amount of solutions must be measured into the same containers, and the procedure of applying them needed to be the same too (time recorded, measurements, temperature, etc. ). The only thing that was different in the next run was that the reactants had to be at a temperature of 10? C above room temperature before mixed.
This was accomplished by placing both the flask with bleach and the graduated cylinder with the cobalt (II) nitrate in a water bath for 10 minutes, and then adding the cobalt (II) nitrate to the flask, then back into the water bath. Hot water was used to increase the temperature, and cold water was used to adjust it. The next run was a similar idea to the previous one, but the reactants were brought down to a temperature 10? C below room temperature using ice. The steps to doing this are similar to the previous ones, but only this one required an addition of 20mL of water to the bleach solution before mixing.
The reason being is so that the overall concentrations are half of their original vales. The run that followed after was also identical, but instead of adding 20mL, 60mL was added. Now the overall concentrations after mixing were one quarter of their original values. The experiment was practically over, but there always had to be cleaning and instructed disposal of chemicals. The product(s) was/were instructed to be disposed in the designated container only for the waste solution. Finally, all the parties that participated in the experiment were obliged to wash their hands thoroughly with soap and water before leaving the laboratory.
Analysis and Results The rate of production of oxygen for each reaction was slightly different. The rate of reaction is determined by the equation; For the control where the reaction to place at room temperature and with bleach with a concentration of 0. 529M, the rate of production of oxygen was 36. 1 mL/minute. In next reaction which took place at a temperature 10? higher than that had a rate of 39. 5 mL/minute. Next was the reaction which took place at 10? below room temperature which resulted in a rate of 26. 8 mL/minute.
In the reaction that 20 mL of distilled water was added to the bleach solution and the temperature was kept constant, the reaction rate dropped to 16. 2 mL/minute. Finally the slowest reaction occurred when 60 mL of distilled water was added to the bleach causing a rate of 10. 8 mL/minute. The rate value changes as the temperature is changed. When the temperature increases by 10? , the rate of the reaction increases by a factor of 0. 12 (12%). This is again changed when the temperature is changed to 10? below room temperature. This results in a rate of production of oxygen, which is decreased by a factor of 0. 5 (25%). When the concentrations were changed so did the rate of reaction. When the concentration was changed to 0. 265M the rate of reaction dropped by a factor of 0. 5 (50%) below the control value. Furthermore when 60mL of water was added to the bleach dropping the concentration too 0. 132M the rate dropped by a factor of 0. 7 (70%). Bleach should never be mixed with any acid based cleaners because it results in the formation of toxic Cl- gas. If bleach is mixed with an acid based cleaner in a small room it will result in a toxic build up of chlorine gas, which can be fatal to anyone spending time in the room.
The equations for these reactions are shown below; Bleach is formed by the action of chlorine gas on sodium hydroxide, NaOH: The equation below represents the reaction of bleach with an acid based cleaner, which gives off chlorine gas Because of this reaction all acid based cleaners have warnings not to be mixed with bleach because it can result in injury or death. If bleach with 10% sodium hypochlorite was used for this experiment instead of bleach with 5. 25% sodium hypochlorite.
The shape of the rate curve for the graph would likely be twice as steep as the graph for the reaction involving bleach with a concentration of 5. 25% sodium hypochlorite. This is because the reaction will finish faster due a concentration that is higher by a factor of two. In this experiment there were possibilities for errors, the main one would be caused by measuring the volume of air at certain times. The reason for this being an error is that at 30 seconds more air will have been produced than is bing measure this is because some oxygen is in the Erlenmeyer flask but still rising to the point at which it is measured.
Also some oxygen is held back because of a kink in the rubber tubing. To improve this experiment I would use a better way of measuring the volume of oxygen produced, either by measuring the air pressure in a container attached to the flask where the reaction was taking place or by using a large tube with a piston inside that would slide along the inside of it showing how much oxygen is evolved during the reaction. Conclusion From the experiment that was carried out it can be concluded that both temperature and concentration effect reaction rates.
The lower the temperature the slower the reaction rate, therefore the higher the temperature the faster the reaction takes place. Likewise the lower the concentration of a solution the slower the reaction and the higher the concentration the faster the reaction is completed. ——————————————– [ 1 ]. Measuring Reaction Rate Using Volume of Gas Produced. ” Experiment 11C. N. p. : SMG Lab, n. d. N. pag. Rpt. in Experiment 11C. N. p. : n. p. , n. d. 154-58. Print. [ 2 ]. DiGiuseppe, et al. Reaction Rates. N. p. : Nelson, 2012. Print. Nelson Education.