Chapter 26

An electrochemical cell A device that converts chemical energy into electrical energy. A Daniell cell is a device that could supply a useful electric current. This device was used almost exclusively to power the early English and American telegraph systems because of its reliable output

 A model of the Daniell cell can be made using a simple laboratory apparatus.  An electric current flows through the wire and light the globe.  This part of the cell is called the external circuit.  The glove converts the electrical energy of the current into light and heat.

The current flows because a chemical reaction is taking place in the cell. Initially there is little indication that a reaction is taking place in the beakers. However if we leave the cell with a light globe connected for several hours a number of changes can be observed: – The zinc metal corrodes – The copper metal becomes covered with a furry brown-black deposit – The blue copper(II) sulfate solutions loses some of its colour These changes provide evidence of a chemical reaction

 If the light glove is replace with a galvanometer, the galvanometer indicates that electrons flow from the zinc electrode, through the wire, and to the copper electrode.  Current flows only if a salt bridge is present.

The reaction in the cell is a redox reaction, as electrons are being produced and consumed. The zinc electrode is eaten away, forming zinc ions in solution: Zn(s) → Zn 2+ (aq) + 2e - The oxidation of the zinc metal releases electrons, which flow through the wire to the copper electrode. Electrons are accepted by copper ions in the solution when the ions collide with the copper electrode: Cu 2+ (aq) + 2e - → Cu(s) The salt bridge forms an essential part of the cell

Cu 2+ (aq) + Zn(s) → Cu(s) + Zn 2+ (aq)  Copper ions act as the oxidant and zinc metal acts as the reductant. What is being oxidised What is being reduced

A galvanic cell is designed so that half reactions occur in two separate compartments of the cell. The oxidants and reductants do not come into direct contact with each other. Electrons can only be transferred by travelling through an external circuit connecting the negative and positive electrodes. The flow of electrons creates an electric circuit. The chemical energy is transformed into electrical energy.

 The Daniell cell can be regarded as consisting of two half cells.  Each half cell contains an electrode in contact with a solution.  The species present in each half cell form a conjugate redox pair

 In general if one member of the conjugate pair in a half cell is a metal, it is usually used as the electrode.  However if no metal is present an inert electrode such as platinum or graphite is used.  In some half cells one of the conjugate pairs may be a gas.

The electrode at which oxidation occurs is called the anode. In galvanic cells the anode, where electrons are produced, is described as the negative terminal The electrode where reduction occurs is called the cathode. In galvanic cells this is where electrons are consumed, it is described as the positive terminal. In the Daniell cell the zinc electrode is the anode and the copper electrode is the cathode.

 The salt bridge connects the circuit  It contains ions that are free to move so that they can balance charges formed in the half cells.  Cations migrate to the cathode and anions migrate to the anode.  This part of the cell is called the internal circuit.

 Why does the zinc metal react with the copper ions rather than the copper metal react with the zinc ions?  They both have a tendency to lose electrons and form ions, but they differ in their tendencies to do so.  Zinc loses its electrons much more readily than copper.  When the Zn 2+ /Zn and Cu 2+ /Cu half cells are connected, electrons flow from the zinc metal in one half cell to the copper ions in the other.

 Zinc is described as a stronger reductant than copper.  Since zinc loses electrons more readily, it follows that the oxidised form of zinc, Zn 2+, will not accept electrons as readily as Cu 2+ ions.  Cu 2+ is described as a stronger oxidant than Zn 2+.

 What are the half cells?  What is being oxidised and what is being reduced?  What is the oxidant?  What is the reductant?  Why is copper oxidised here but it is reduced in the Daniell Cell?

 Remember from last year that the electrochemical series lists the half cell equations in order of their tendency to occur as reduction reactions.

 Is a list of various half cells.  Is only valid under conditions in which it was determined.  The one in your book applies at:  A temperature of 25°C  Pressure of 1 atm  1 M concentrations of solutions  These are standard laboratory conditions (SLC)  We can use the electrochemical series to predict what will happen when two specific half cells are combined to from a cell.

 Use the electrochemical series to predict the behaviour of a cell composed of Ag + /Ag and Fe 2+ /Fe half cells.

 Page 419  Questions 1 and 2

 A galvanic cell is to be constructed using magnesium metal, silver metal and solutions of magnesium nitrate and silver nitrate.  Draw and label a diagram with  The direction of the flow of electrons in the external circuit  The metal electrodes and the electrolytes in the half cells  The appropriate electrolyte in the salt bridge  The polarity of each electrode  Write the reaction occurring at the anode.

 Sometimes called the emf.  It is the electrical pressure between two half cells.  It is measured in volts. It measures the amount of energy supplied by a fixed amount of charge flowing from a galvanic cell.  A Daniell cell has a potential difference of about 1 volt.  We can get a rough indication of the potential difference from the electrochemical series

 We can use the E° values from the electrochemical series.  These are standard half-cell potentials and give a numerical measure of the tendency of a half cell reaction to occur as a reduction reaction.  The hydrogen half cell is the reference cell and has an E° value arbitrarily assigned as zero.  Everything else is compared to this cell.

 The potential difference of a cell at standard conditions is the difference between the E° values of its two half cells. Cell potential difference = higher half cell E° – lower half cell E°  For example Ag + /Ag and Fe 2+ /Fe half cells. Cell potential difference = E°(Ag + /Ag) – E°(Fe 2+ /Fe) = 0.80 – (-0.41) = 1.21 V

 Page 423  Question 3  Predict what will happen, if anything, if some iron filings are added to a solution of tin(II) chloride.

 In a galvanic cell, the higher half reaction in the electrochemical series occurs in the forward direction (reduction).  The lower reaction occurs in the reverse direction (oxidation)  For a reaction to occur, a chemical on the left of the electrochemical series must react with a chemical on the right that is lower in the series.

 Consider the following equations that appear in the order shown in the electrochemical series:  Br 2 (aq) + 2e - 2Br - (aq)  Ni 2+ (aq) + 2e - Ni(s)  Mg 2+ (aq) + 2e - Mg(s)  Predict the effect of mixing  Br 2 (aq) and Mg 2+ (aq)  Mg 2+ (aq) and Ni(s)  Ni 2+ (aq) and Mg(s)

 The standard half cell potentials are measured under standard conditions.  The half cell potentials vary if you move away from the standard conditions.  The order of the half cell reactions can change and the standard half cell potentials will not be relevant.  The electrochemical series gives no information about the rate at which reactions occur.  According to the electrochemical series hydrogen peroxide would form water and oxygen gas but this reaction takes years to occur unless a catalyst is present.

 Page 425  Question 4  Page 436  Question 6 – 9, 12, 14