Lab Report on Rate of Reaction PDF

Preview

Summary

An Experiment to Determine the Effect of Temperature on Rate of Reaction Between Sodium Thiosulfate and Hydrochloric Acid...

Description

FND04: CHEMISTRY LAB REPORT

168406

An Experiment to Determine the Effect of Temperature on Rate of Reaction Between Sodium Thiosulfate and Hydrochloric Acid Introduction Chemical kinetics is defined as the study of chemical reactions in terms of their reaction rates and the various factors that affect them. The rate at which the reaction occurs can be altered by changing the set of conditions such as: temperature, concentration, pressure, surface area and the use of catalysts (Lister and Renshaw, 2008). Simple collision theory highlights the way in which these factors affect the reaction rate on a molecular level. In order for a reaction to occur, the particles must collide with sufficient energy to break bonds and must also be in the correct orientation. The collision must take place between the area of the molecule that is going to react together especially with unsymmetrical species as repulsion may cause molecules to bounce off each other. Moreover, even with the correct orientation a reaction cannot occur unless the reactant particles collide with the minimum energy called the activation energy (Ea). The activation energy is the barrier in which a reaction can take place and so only those collisions which have energies equal to or greater than the activation energy can result in a reaction. Therefore, the rate of rate of reaction is dependent upon the rate of successful collisions between the reactant particles (The Editors of Encyclopædia Britannica, 1998). Altering certain factors affecting reaction rates will determine the speed of a chemical reaction. For instance, by increasing the temperature of the reaction the kinetic energy of the particles also increases and so a greater proportion of colliding particles will have the sufficient energy to overcome the activation energy and react. The Maxwell-Boltzmann distribution graph displays the distribution of velocities of the reactant particles’ kinetic energy which gives an understanding of the effect of the factors of rate of reaction. As shown in Figure 1, the shaded region for T2 is a lot Figure 1: Maxwell-Boltzmann distribution for the particles for a reaction mixture at two different larger than T1 which infers that there is a greater proportion of temperatures (Fullick and McDuell, 2008). particles which have overcome the activation energy at a higher temperature (Fullick and McDuell, 2008). For many reactions that occur around room temperature (25˚c), the rate of reaction approximately doubles for every 10˚c rise in temperature (Curtis, Hunt, and Hill, 2015). The rate of reaction differs depending on the substances involved for instance, ionic precipitation reactions are rapid whereas corrosion processes such as the rusting of iron occur slowly (Curtis, Hunt, and Hill, 2015). This suggests that the nature and strength of the bonds of the reactants heavily impacts the reaction rate. Furthermore, there are other factors which affect the rate of reaction such as the surface area of solid reactants, the larger the surface area of the solid, the more sites for reaction meaning more particles are available to collide with. Increasing pressure and concentration have similar effects in which there are more particles in a given volume and so there is a higher collision frequency. Using a catalyst increases the rate of reaction by following a different pathway for the reaction to follow. These factors all contribute to increase the rate of reaction by increasing the effective collision of particles that occur in a reaction (Lister and Renshaw, 2008).

FND04: CHEMISTRY LAB REPORT

168406

The rate of reaction between sodium thiosulfate and dilute hydrochloric acid can be measured by the time it takes for the sulfur precipitate to be formed, as shown in the equation below: 2HCl(aq) + Na2S2O3(aq) 

2NaCl(aq) +

SO2(aq) + S(s) + H2O(l)

Hydrochloric Acid + Sodium Thiosulfate  Sodium Chloride + Sulfur Dioxide + Sulfur + Water

The result of this equation forms sulfur dioxide which is a toxic gas at room temperature but is very soluble in water whereas sulfur is insoluble. This explains the physical appearance of the reaction in which the solution first becomes cloudy and gradually turns opaque due to the precipitate formed by the solid sulfur. The time taken for the solution to turn opaque is used to determine the rate of reaction by using the following equation: Rate of reaction = 1 / time taken for the solution to turn opaque In order for the effect of temperature on rate reactions to be investigated, all the other factors must be kept constant. Hypothesis As the temperature of the mixture increases, the rate of reaction will increase. This is due to the increase in kinetic energy of the reactant molecules, hydrochloric acid and sodium thiosulfate and will have more effective collisions. For every 10˚c rise in temperature, the rate of reaction roughly doubles which suggests that the graph will be of an exponential curve shape. Therefore, as the rate of reaction increases along with the temperature, the time taken for the cross to disappear decreases as the formation of the precipitate will be faster. Method The sodium thiosulfate solution was measured into a conical flask and placed in a water bath of a selected temperature. The conical flask was left in the water bath for around five minutes and monitored by a thermometer to determine whether the solution has reached the temperature of the water bath. Following this, an X mark was drawn on a ceramic tile and placed underneath the conical flask. This X mark is important in the experiment to determine the rate of reaction by the time it takes for the mark to disappear. The hydrochloric acid is added into the sodium thiosulfate solution in the conical flask and timer is set immediately. It was vital to start the timer as soon as the acid was added to accurately measure the time it would take for the reaction to occur. The flask was gently swirled in circular motions until the X mark below the conical flask was no longer visible due to the clear solution turning opaque. The contents of the conical flask were then poured into a sodium carbonate solution in a fume cupboard due to the production of sulfur dioxide in the reaction which is highly toxic and causes breathing difficulties. This method was then repeated for various temperatures – 25˚c, 35˚c, 45˚c and 55˚c.

FND04: CHEMISTRY LAB REPORT

168406

Results Table 1: Quantitative data collected from the rate of reaction in various temperature Temperature of the Time taken for the Experiment mixture in the flask / cross to disappear / Rate of reaction / s-1 number ˚c s 25 155 0.00645 35 91 0.0110 1 45 52 0.0192 55 38 0.0263 25 134 0.00746 35 73 0.0137 2 45 47 0.0213 55 33 0.0303 25 170 0.00588 35 105 0.00952 3 45 59 0.0169 55 41 0.0244

Qualitative Observation: The sodium thiosulfate and hydrochloric acid solutions were colourless. However, once the hydrochloric acid was added to the sodium thiosulfate in the conical flask, the solution gradually changed from colourless into a cloudy cream colour. Eventually, the X mark slowly disappeared as time went on due to the cloudiness of the solution.

Graph 1 showing the rate of reaction in various temperatures 0.035

Rate of reaction / s-1

0.03 0.025 Experiment 1 0.02

Experiment 2

0.015

Experiment 3 Expon. (Experiment 1)

0.01

Expon. (Experiment 2) Expon. (Experiment 3)

0.005 0 0

10

20

30

Temperature / ˚c

40

50

60

FND04: CHEMISTRY LAB REPORT

168406

Graph 2 showing the mean rate of raction in various temperatures 0.14

Rate of reaction / s-1

0.12 0.1 0.08 0.06 0.04 0.02 0 0

10

20

30

40

50

60

Temperature / ˚c

Conclusion Based on experimental data, the rate of reaction increases as the temperature increases which supports my hypothesis to an extent. Graph 1 indicates that experiment 2 shows the most reliable result in terms of fitting the trend stated in my hypothesis in which the rate of reaction roughly doubles for every 10˚c increase in temperature. In comparison to my own experiment (experiment 1) and experiment 3 in which the rate is not as consistent. Furthermore, Graph 2 highlights the limitations of a small set of data from other individuals due to human errors. Although graph 2 also follows the trend in which the rate of reaction is dependent of the temperature rising, the doubling of the rate for every 10˚c is not so clear which contrasts my hypothesis. The results show that the dependent variable, in this case the rate of reaction is proportional to the independent variable, the temperature in this investigation. This is because the collision theory states that an increase in temperature causes a rise in kinetic energy thus a higher collision frequency will cause a reaction to occur faster. The higher the temperature, the more particles overcome the activation energy. In this investigation, the reaction between hydrochloric acid and sodium thiosulfate caused the X mark to disappear due to the formation of a precipitate of sulfur. The colourless solutions reacted together and shifted to cloudy followed by being opaque. This reaction can be summarised by the following ionic equation: S2O32- (aq) + 2H+ (aq)  S (s) + SO2 (aq) + H2O (l) It was vital for there to be controlled variables due to the fact that the various factors such as the volume and concentration of the solution, would have influenced the rate of reaction.

FND04: CHEMISTRY LAB REPORT

168406

Overall, the investigation of the effect of temperature on the rate of reaction was successful in a sense that the relationship between the two variables coincided and my results supported my hypothesis in showing that as the temperature rose, so did the rate of reaction. However, some minor faults to the experiment due to systematic errors would have impacted my results and so an exponential shaped graph was not fully achieved. Evaluation Although the results of this experiment followed the predicted trend, there were still some errors made in my method and certainly some improvements to be made. Due to time constraints, I only had the chance to do one experiment and so to make my graph more reliable I gathered the results of other individuals (experiment 1 and 2 of graph 1) to compare with my set of results. This was not particularly ideal as their method of this experiment may have differed from my own and so the rate of reaction would have been affected by several factors such as the size of the X mark and the method of swirling the conical flask. Also, different people would have different interpretations of the mark to be completely disappeared. In future, I would do more trials by myself to get more reliable results and eliminate random errors. It was difficult to accurately measure the time it took for the X mark to disappear due to various factors. For instance, the X mark that I drew was quite large therefore it took much longer for this mark to disappear as the solution gradually shifted from colourless to cloudy to opaque however this took some time for it to fill the base of the conical flask. In future, the X mark on the ceramic tile could have been drawn to a suitable size to fit the centre of the base of the conical flask. Furthermore, the swirling of the conical flask in circular motions was not entirely consistent in the duration of the experiment. Therefore, this factor may have affected the time it took for the X mark to disappear and in turn impacted the rate of reaction in the results. This could have been avoided by using a magnetic stirrer which would stir the mixture of solutions at a constant speed which would make the results more accurate. It was sometimes difficult to detect the disappearance of the X mark due to visibility problems. To make the results a lot more accurate and without human error, a light sensor would have been ideal in this investigation in order to measure the light that passes through the solution. Therefore, once the precipitate has formed and the solution has turned cloudy, light would be impenetrable which would be picked up by the sensor. In turn, the measurement of time taken for the X to disappear and the solution to turn fully opaque would be a lot more precise. Due to the quick pace that I was working at, the conical flask and measuring cylinders may not have been washed completely leaving traces of the previous solution behind. This would result in inaccurate measurements of the reactants which in turn would affect the overall rate of reaction. This could have been improved by being patient and more mindful of the equipment that I was using to achieve accurate results and maintaining a consistent method during each trial.

FND04: CHEMISTRY LAB REPORT

168406

References Curtis, G., Hunt, A. and Hill, G. (2015) Edexcel A Level Chemistry Student, Book 1. London, United Kingdom: Hodder Education. The Editors of Encyclopædia Britannica (1998) ‘Collision Theory’, in Encyclopædia Britannica. Available at: https://www.britannica.com/science/collision-theory-chemistry (Accessed: 14 December 2016). Fullick, A. and McDuell, B. (2008) Edexcel AS Chemistry Student Book. Harlow: Longman. Lister, T. and Renshaw, J. (2008) AQA AS Chemistry. Cheltenham, United Kingdom: Nelson Thornes....