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Electrochemical Cells

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Electrochemical Cells Tam Paul-Worika Contents Aim of the Investigation.....................................................3 Background.....................................................................3 * Redox Reactions * Electrochemical Cells * Salt Bridge * Nernst Equation Plan.................................................................................5 * Equipment * Risk Assessment * Instructions * Making the solutions Analysis...........................................................................7 Evaluation......................................................................17 Aim of the Investigation In my investigation, I will observe how changing the concentration of a solution can affect the electrode potential in an electrochemical cell. I will also see if there is a trend when the concentration is changed. Background Redox Reactions Redox reactions are reactions where the oxidation states of the atoms change[1]. The atoms are either oxidised or reduced, depending on if they lose or gain electrons. Electrochemical Cells Electrochemical cells (also known as Galvanic cells or Voltaic cells[2]), are devices that cause a current from chemical reactions. These reactions occur through redox. An electrochemical cell is set up so that the electrons lost from one of the reagents can travel to another reagent. This creates a voltage, which is also known as the electric potential difference. This voltage can be read if a high-resistance voltmeter is put into the circuit. Electrochemical cells are made from two half-cells, which have their own standard electrode potential (E0), which is measured in volts. This is measured by placing it with a standard hydrogen electrode. ...read more.


May cause long-term adverse effects. Copper (II) Sulphate: Dissolve in at least 8 litres of water for 40g before disposal. Zinc Sulphate Zinc Sulphate: Dissolve in at least 4.05 litres of water for 40.5g before disposal. Potassium Nitrate Oxidising Agent. Contact with a combustible material may cause fire. For 100g, dissolve in 10 litres of water before disposal. Zinc Dust (from cleaning the zinc strips) Flammable. Contact with water releases hydrogen gas Destroy by adding slowly, with stirring, to excess dilute acid Instructions 1. Make a 1mol dm-3 solution of copper (II) sulphate and zinc sulphate. 2. Pour 50cm3 of each into 100cm3 beakers. Using a folded piece of filter paper soaked in saturated potassium nitrate for the salt bridge, and a voltmeter (as shown in fig. 3), measure the electrode potential between the two solutions. Measure the temperature of the solutions at the time of the reaction. 3. Repeat number 2 for all of the different concentrations (1, 0.5, 0.25, 0.125, 0.625mol dm-3) 4. Keeping the copper (II) sulphate at 1mol dm-3, change the concentration of the zinc sulphate Making the Solutions All of the solutions that I am going to use in my investigation have to be made from the solid. ...read more.


sulphate decreases. The half-cell potential of copper (II) sulphate also increases when the concentration decreases. This happens because as the concentration decreases, the number of electrons in the solution is less than usual. Because of this, the solution is even more likely to receive more electrons, making it more positive. The half-cell potential of zinc sulphate also increases because it is less likely to give away electrons than usual. In the graphs, you can see that when you compare the electrode potentials for each concentration of copper (II) sulphate, all of the lines are almost parallel to each other. This shows that if, for example, you have 1mol dm-3 of copper (II) sulphate and 0.5mol dm-3 of zinc sulphate, you can tell that the electrode potential is 1.084V. If you change the concentration of copper (II) sulphate to 0.5mol dm-3, you would have to add 0.020V to it, and then if you had 0.25mol dm-3 of copper (II) sulphate, you add 0.021V. For each one, you add around 0.020V, so you can work out the approximate electrode potential. Evaluation Measurement Errors Beakers: 0.05ml/50ml x 100 = 0.1% 0.05ml/200ml x 100 = 0.025% 0.05ml/250ml x 100 = 0.02% Measuring Cylinders: 0.05ml/50ml x 100 = 0.1% Thermometer: � 0.5�C Voltmeter: � 0.0005V Volumetric Flask: 0.5ml/250ml x 100 = 0.2% Burette: 0.05ml/25ml x 100 = 0.2% Scales: 0.0005g/40g x 100 = 1/800% 0.0005g/40. ...read more.

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