Lab Report on Oxidation and Reduction PDF

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REDOX lab report...

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FND04: FURTHER CHEMISTRY FOR BIOSCIENCES - LAB REPORT Oxidation and Reduction (REDOX) Reactions Introduction Oxidation-reduction reactions, otherwise known as redox reactions, are fundamentally a process that involves a transfer of electrons. Oxidation can be defined as loss of electrons and reduction as gain of electrons. Therefore, oxidation results in an increase in oxidation number and reduction a decrease in oxidation number (Gent and Ritchie, 2007). In a redox reaction, the substance that is reduced is called an oxidising agent as it gains electrons from the substance that is oxidised. In contrast, a reducing agent loses electrons so it itself is oxidised whilst reducing another substance (Gent and Ritchie, 2008). Although oxidation and reduction reactions can be viewed separately through half-equations, reduction and oxidation occur simultaneously forming a whole reaction (Haustein, 2014). The total number of electrons lost must equal the total electrons gained in a redox reaction in order to maintain neutrality. The half equations can be combined to construct a balanced full net ionic equation. For instance, in the reaction of iron (III) chloride and tin (II) chloride, the iron is being reduced and the tin is being oxidised. The chloride ions are unaffected and are known as spectator ions. Fe3+ + e-  Fe2+ Sn2+  Sn4+ + 2e-

Reduction Oxidation

The following half equations can be combined by balancing them through multiplying the coefficients in the reduction equation by 2. Followed by adding the balanced half equations together. Sn2+  Sn4+ + 2e2Fe3+ + 2e-  2Fe2+ Sn2+ + 2Fe3+ + 2e-  Sn4+ + 2Fe2+ + 2eSn2+ + 2Fe3+  Sn4+ + 2Fe2+

(Blackman et al., 2015)

Substances that have a large affinity for electrons tend to be good oxidising agents such as potassium manganate and the halogens. In comparison, metals are usually good reducing agents as they have a strong tendency to lose electrons (Gent and Ritchie, 2007). Redox agents are responsible for the visible colour changes as well as changes in property in some instances. This is because a colour change indicates that a chemical reaction has taken place. Furthermore, the products of a redox reaction can then be confirmed by adding certain reagents to test for it. For instance, iodine is a reducing agent which will turn blue/black in the presence of starch. Redox reactions have various applications in chemistry such as reducing alcohol to aldehyde which is essential in organic synthesis. Furthermore, reduction and oxidation reactions are vital to some of the basic functions of everyday life including photosynthesis, respiration, combustion and rusting (Haustein, 2014).

Aim To observe the redox reactions of common oxidising and reducing reagents as well as identifying which of the chemical substances are being oxidised and reduced by constructing equations for each of the 6 reactions. Method This practical has been split into 6 parts. Firstly, 5cm3 of iron (II) sulphate solution and 2cm3 of dilute sulphuric acid were added into a test tube. Followed by a dropwise addition of 2cm3 of potassium (VII) manganate. Observations of the addition of potassium (VII) manganate were then noted as the manganate ions (MnO4-) acted as oxidising agents thus oxidising the iron (II) ions (Fe2+). After this, the solution was divided into two separate test tubes. To one, a few drops of potassium hexacyanoferrate (II) solution was added and to the other, potassium thiocyanate solution. These reagents confirmed the presence of the products formed by the change in colour of the two solutions. The second experiment follows, 5cm3 of potassium iodide solution and 2cm3 of bromine water were put together in a test tube. The product of this reaction was tested by adding a few drops of starch solution to test for iodine ions (I-). An observation of this reaction was then noted. Following this, 2cm3 of iron (III) chloride solution was added into a test tube followed by 2cm3 of potassium iodide. The resulting solution was split equally into two test tubes. To one, starch solution was added to test for I- and to the other, sodium thiosulphate was added until no further colour change occurred followed by adding a few drops of potassium ferricyanide (potassium hexacyanoferrate (III)) solution to confirm the presence of iron (II) ions. The colour changes of these reactions were observed. Experiment four consisted of adding 2cm3 sodium sulphite solution into a test tube as well as 2cm3 of dilute sulphuric acid and 1cm3 of potassium dichromate. This resulting solution was then warmed gently by placing the test tube into a beaker of warm water whilst swirling the test tube in order to speed up the reaction. The dichromate ions (Cr2O72-) act as oxidising agents and so results in a colour change which was observed. Experiment five consisted of adding 5cm3 of concentrated hydrochloric acid in a boiling tube and 5cm3 of potassium manganate (VII) solution. The resulting solution was also warmed gently by placing the test tube in a beaker of warm water. An observation of the colour change was made followed by a test of the gas evolved. A damp universal indicator paper was placed over the top of the boiling tube and the colour of the paper indicated the pH of the gas. Lastly, experiment six comprised of adding 2cm3 of sodium sulphite solution as well as 2cm3 of chlorine water. The products of the reaction was then tested by adding a few drops of barium chloride solution and an observation was made.

Results

Experimen t

Observation When dropwise addition of the potassium manganate (VII) solution was added to the iron (II) sulphate and dil. sulphuric acid solution, the mixture remained colourless.

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When potassium hexacyanoferrate (II) solution was added, the solution turned from colourless to deep blue. When potassium thiocyanate solution was added, the solution turned from colourless to a red-brown colour. When starch solution was added to a solution of potassium iodide and bromine water, the solution turned from orange to black. Adding starch to the solution of iron (III) chloride and potassium iodide turned the solution from brown to black. Adding sodium thiosulphate to the solution turned the solution black then went to an orange-brown solution. Adding a few drops of potassium ferricyanide (potassium hexacyanoferrate (III)) after no further colour changes occurred, the solution turned a dark green colour. When potassium dichromate was added to a warmed solution of sodium sulphite and dilute sulphuric acid, there was no observable reactions as the solution remained an orange colour throughout the experiment. Adding potassium manganate (VII) solution to the concentrated hydrochloric acid caused the solution to turn brown and when warmed gently, turned slightly clearer and appeared light brown. When testing the gas evolved with a damp universal indicator paper, the red indicated a pH of 1. Testing the sodium sulphite and chlorine water solution with barium chloride caused the solution to turn from colourless to a cloudy white precipitate.

Conclusion In the first experiment, the solution was initially purple due to the potassium manganate (VII) solution but then turned colourless indicating that the MnO4- ions had been reduced to Mn2+ ions. Supposedly, a purple layer should have formed after reaching a certain point but the result does not show this. Adding the potassium thiocyanate solution resulted in the solution turning red-brown indicating the presence of Fe3+. The expected colour for this was supposed to be blood red. The presence of Fe3+ was also confirmed by the addition of potassium hexacyanoferrate (III) solution as the solution turned from clear to a deep blue colour. The expected outcome of this test was meant to give a Prussian blue precipitate. MnO4-(aq) + 8H+(aq) + 5e- Mn2+(aq) + RED 4H2O(l) OX Fe2+(aq) Fe3+(aq) + eMnO4-(aq) + 8H+(aq) + 5Fe2+ Mn2+(aq) + 4H2O(l) + 5Fe3+ (aq) In this case, the MnO4- ions are being reduced therefore acting as the oxidising agent in this reaction and Fe 2+ as the reducing agent. Experiment 2 showed the potassium iodide and bromine water solution turned from orange to black when tested with starch solution. The colour change implies that iodine is present as a product of the reaction. Therefore, Br2 acts as the oxidising agent and I- act as the reducing agent. The result of this experiment follows the expected outcome as iodine turns starch blue/black. Br2 (aq) + 2e- 2Br-(aq) 2I-(aq) i2 (aq) + 2e-

RED OX

Br2 (aq) + 2I-(aq) I2 (aq) + 2Br-(aq) Experiment 3 showed a colour change when starch solution was added from a brown solution to black. The expected result for this was meant to be a dark purple solution instead of black. This implies that iodine is present as a product which is true in this case as I- is oxidised to I2 by the oxidising agent Fe3+. Furthermore, when sodium thiosulphate solution and potassium ferricyanide solution was added, the solution turned dark green indicating the presence of Fe2+. The expected outcome of the addition of sodium thiosulphate was meant to be a dark purple solution but goes back to the original brown solution over time. The addition of potassium ferricyanide was meant to form a dark turquoise colour change. Fe3+(aq) + e- Fe2+(aq) 2I-(aq) i2 (aq) + 2ef 2Fe3+ (aq) + 2I-(aq) 2Fe2+ (aq) + I2 (aq)

RED OX

Experiment 4 did not react according to what was expected. The solution should have turned from orange to green to indicate the reduction of Cr2O72- to Cr3+. SO32- is the reducing agent in this experiment as it reduces the dichromate ions and therefore the dichromate ions act as the oxidising agent. RED Cr2O72-(aq) + 14H+(aq) + 6e- 2Cr3+(aq) + 7H2O(l) SO32-(aq) + H2O(l) SO42-(aq) + 2H+(aq) + 2e-OX Cr2O72-(aq) + 8H+(aq) + 3SO32-(aq) 2Cr3+(aq) + 4H2O(l) + 3SO42-(aq) Experiment 5 showed that the solution of concentrated hydrochloric acid and potassium manganate (VII) solution turned brown which was not the expected observable result. The addition of the potassium manganate (VII) solution should have turned the colourless solution red. Therefore, Cl- act as the reducing agent as it reduces MnO4- to Mn2+ causing the colour change. Furthermore, the universal indicator paper showed a very acidic pH of 1 which was a bit unexpected as the brown precipitate produced was manganese oxide (MnO2) indicating that insufficient acid was added. RED MnO4-(aq) + 8H+(aq) + 5e- Mn2+(aq) + 4H2O(l) OX 2Cl-(aq) Cl2 (aq) + 2ef 2MnO4-(aq) + 10Cl-(aq) + 16H+(aq) 2Mn2+(aq) + 8H2O(l) + 5Cl2 (aq) Experiment 6 coincided with the expected result of the reaction. The barium chloride solution turned the resulting solution from colourless to a cloudy white precipitate which is a common test for sulphate ions. Cl2 is responsible for reducing SO32- to SO42- in order for the product, SO42- to be tested. RED Cl2 (aq) + 2e- 2Cl(aq) + H2O(l) SO42-(aq) + 2H+(aq) OX SO32-(aq) Cl2 (aq)++2eSO32-(aq) + H2O(l) 2Cl-(aq) + SO42-(aq) + 2H+ (aq) Evaluation Most of my results coincide with the expected outcomes of the colour change tests in each redox reaction. However, there was a few experiments in which the anticipated reaction did not occur. For instance, experiment four did not work in the practical as no visible change was observed. Various sources of error may explain the reasoning behind this such as insufficient cleaning of the test tubes in between experiments therefore chemical residue may be left in the tubes affecting the products of the reaction taking place. Furthermore, this experiment required heat in order for the reaction to occur instantly on the basis of chemical kinetics as potassium dichromate is a less powerful oxidising agent than potassium manganate (VII). Another source of error could be accuracy of measuring out the supplied reagents. Similarly, experiment five had conflicting results as the pH test indicated the solution to be very acidic however the brown precipitate implies that insufficient acid was added producing MnO2. This could also be explained by inaccurate measurements of the

reagents and could be improved by carefully measuring out the chemicals using appropriate equipment such as pipettes. For future practical’s, more care should be given to the measurements of the chemicals given to provide accurate results. Moreover, the qualitative analysis was not as precise in comparison to the expected colour changes. For instance, the colour change in experiment three was meant to turn a dark turquoise colour with the addition of potassium ferricyanide however my results describe the colour as dark green. Likewise, experiment one was meant to be Prussian blue instead of deep blue. The preciseness of descriptive observations would have greatly improved the accuracy and reliability of the practical. However, the descriptive analysis is difficult to completely match the expected as different people would have different ways of describing a shade of colour. References Blackman, A., Bridgeman, A., Laurie, G., Southam, D., Thompson, C. and Williamson, N. (2015). Chemistry: Core concepts. 1st ed. Gent, D. and Ritchie, R. (2007). OCR Chemistry AS. 1st ed. Oxford: Heinemann. Gent, D. and Ritchie, R. (2008). OCR A2 Chemistry. 1st ed. Oxford: Heinemann. Haustein, C. (2014). Oxidation-reduction reactions. [online] Gale Science in Context. Available at: http://link.galegroup.com/apps/doc/CV2644031629/SCIC? u=albertak12&xid=5e56703a [Accessed 22 Mar. 2017]....