Metal and Non-metal Redox Reactions Experiment

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Abstract This experiment aimed to investigate Redox reaction and hence determine which elements were reactive. For this experiment, the practical was performed separately, metal versus metal redox reactions and non-metal versus non-metal reactions. The metals in question were zinc, lead and copper which were reacted separately against each other’s nitrate salts after which observations were made. Similarly, the non-metals in question were aqueous solutions of chlorine, iodine and bromine which were reacted against each other’s sodium salts.

The experimental results revealed that redox reactions took place as reflected by colour changes. Also, it was established that for the metals, the reactivity series increases from lead, copper, to zinc. On the other hand, for non-metals, the reactivity increases from chlorine, bromine, to iodine. Introduction Also known as oxidation-reduction reactions, Redox reactions are synonymous with acid-base reactions (Albert & Serjeant, 1971). Basically, redox reactions represent a family of reactions involving the exchange of electrons between species involved (Reichardt, 2003). Akin to acid-base reactions, these reactions are a paired type i.e., a two-way traffic reaction that happens concurrently. Noteworthy, while reduction occurs courtesy of electron gain, oxidation is the directly opposite (Department of Chemistry, 2008). Objective This experiment aimed to investigate Redox reaction and hence determine which elements were reactive. Procedure As regards this experiment, metals of zinc, copper, and lead were reacted one at a time with 1.5 ml. of 0.1 M. nitrate salts of zinc, copper, and lead. This was left to settle for 5-10 minutes, an observation made, and, later, tabulated for discussion. Similarly, for non-metals, 1.5 ml. of 0.1 M. sodium salts of bromide, iodide and chlorine in separate test tubes were reacted one at a time with bromine, iodine and chlorine gases but separately. This was left to settle for 15 seconds after which the results were tabulated for analysis. Of note, before this, 1.5 ml. dichloromethane was reacted separately with aqueous solutions of chlorine, bromine and iodine after which results were tabled for analysis 15 seconds later. Results Table 1: observation of metal versus metal reactions. Zn (s) Cu(s) Pb(s) Zn (No3) Same colour Same colour Same colour Cu (No3) Black Same colour Black Pb (No3) Black Same colour Same colour Table 2: observation of non-metal verses non metal reactions (dichloromethane vs. non metal after 10 min). e

Cl2 (aq) Br2 (aq) I2 (aq) Time After 10 min After 10 min After 10 min Colour of organic phase (CH2Cl) Same colour Brown Colour Pink Colour Table 3: observation of non metal verses non metal reactions (organic salts of chlorine vs. non metal). NaBr NaI Colour Remained Clear Remained Clear Here stayed clear but there was no change in colour Here stayed clear but there was no change in colour Table 4: observation of non metal verses non metal reactions (inorganic salts of chlorine vs. none metals). Cl2 Br2 I2 NaCl X Clear to yellow Clear to pink NaBr Clear to Pink NaI Stayed clear Clear to Brown X Discussions and Analysis From the results tabled above, zinc metal was oxidized by both metals. Copper was oxidized by lead metal which was not oxidized at all. As portrayed by table 1 above, a reaction between zinc metal verses either salt of copper and lead metals effected colour changes to black. However, the reverse of the reactions experiences no colour change. The same happens to copper verses lead nitrate though because the two have the same colours it is difficult to notice (Reichardt, 2003). The equations of ‘half reactions’ for metals where oxidation took place in decreasing order are given as below: Zn?Zn2+ + 2e Zn?Zn2+ + 2e Cu?Cu2+ + 2e (Reichardt, 2003) Corresponding reduction ‘half reactions’ for the above reactions are shown below: Pb + 2e?Pb2+ Cu + 2e?Cu2+ Pb + 2e?Pb2+ (Reichardt, 2003) Balanced total reactions for the redox reactions which took place are as show: Zn +Pb2+ ?Zn2++Pb Zn +Cu2+?Zn2+ +Cu Cu+Pb2+? Cu2++Pb Zn2+ +Cu? Zn2+ +Cu (no change) Zn2++Pb? Zn2++Pb (no change) (Reichardt, 2003) For the halide ions, iodine gas was oxidized by both halides. Bromine gas was oxidized by chloride ions which were not ionized at all. As such, as portrayed by table 4 above, reactions between iodine and sodium salts of bromine and chlorine effect no color change remains clear pink, the colour of iodine gas. The reverse reactions give the same colour.

The ‘half reactions’ for the halides where oxidation took place in decreasing order are shown as below: 2I–?I2+2e 2Br–?Br2+2e 2I–?I2+2e (Reichardt, 2003) The corresponding ‘half reactions’ for the above oxidation equations are: Cl2+2e?2Cl– Cl2+2e?2Cl– Br2+2e?2Br– (Reichardt, 2003) Balanced total reactions for the Redox reactions which took place are as show: 2I–+ Cl2? I2+2Cl– 2Br–+ Cl2? Br2+2Cl– 2I–+ Br2? I2+2Br– 2Cl–+I2?2Cl–+I2 (no change) 2Cl–+Br2?2Cl–+Br2 (no change) Conclusion The objective aimed to investigate Redox reactions and hence determine the trend of the reactivity series. This objective was achieved since it was established, through colour changes that indeed redox reactions took place. Also, it was established that for the metals, the reactivity series increases from lead, copper, to zinc. On the other hand, for non-metals, the reactivity increases from chlorine, bromine, to iodine. References Albert, A., & Serjeant, E.P. (1971). The Determination of Ionization Constants: A Laboratory Manual. Kansas City: Chapman & Hall. Atkins, P.W., & Jones, L. (2008). Chemical Principles: The Quest for Insight (4th Ed.). Vatican City State: W.H. Freeman. Department of Chemistry. (2008). Redox Reactions. Web. Hawthorne, A., & Thorngate, J. (1979). Application of Spectroscopy. Birmingham City, UK: University of Alabama.

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Hulanicki, A. (1987). Reactions of acids and bases in analytical chemistry. New York City: McGraw-Hill Companies Inc. Kenkel, J. (1994). Analytical Chemistry for Technicians. Boca Raton, US: Lewis Publishers. Perrin, D.D., Dempsey, B., & Serjeant, E.P. (1981). pKa Prediction for Organic Acids and Bases. Kansas City: Chapman & Hall. Reichardt, C. (2003). Solvents and Solvent Effects in Organic Chemistry: Solvent Effects on the Position of Homogeneous Chemical Equilibria. New York City: John Wiley & Sons, Inc. Scorpio, R. (2000). Fundamentals of Acids, Bases, Buffers & Their Application to Biochemical Systems. Birmingham City, UK: University of Alabama. Sigurds, S. (1986). Applications Of UV-Visible Number UV-31 Derivative Spectrophotometry. Steinhauserstrasse, Switzerland: Peter Lang Publishing Group. Skoog, D., West, D., & Holler, F. (1992). Fudamentals of Analytical Chemistry. Fort Worth, US: Saunders College Publishing. Skoog, D.A.; West, D.M.; Holler, J.F.; Crouch, S.R. (2004). Fundamentals of Analytical Chemistry (8th ed.). Salt Lake City: Thomson Brooks/Cole Publishers. Steiner, J., Termonia, Y., & Deltour, J. (1972); Analitical Chemistry. Geneva: ILO Publications. Szalay, L. (2008). Atomic Absorption Spectrophotometry (AAS). Budapest, Hungary: Petrik Lajos Publications. Vandenbelt, J., & Henrich, C. (1953). Application Spectroscopy. Geneva: ILO Publications.

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