(c) doc b

Introduction to simple cells and batteries - how do they work?

Doc Brown's Chemistry KS4 science–chemistry GCSE/IGCSE/O level/A Level Revision

ELECTROCHEMISTRY revision notes on electrolysis, cells, experimental methods, apparatus, batteries, fuel cells and industrial applications of electrolysis

10. Simple Cells and Batteries

See also 11. Fuel Cells e.g. the hydrogen - oxygen fuel cell

How a simple cell can be used as a battery is explained, using the different reactivities of two metal strips. These revision notes on how simple cells and batteries work should prove useful for the new AQA chemistry, Edexcel chemistry & OCR chemistry GCSE (9–1, 9-5 & 5-1) science courses.

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A simple voltaic cell from two metal strips dipped in an acid or salt solution (c) doc b10. Simple Cells and batteries

  • In electrolysis, electrical energy is taken in (endothermic) to enforce the oxidation and reduction to produce the products at the electrodes.

  • The chemistry of simple voltaic cells or batteries is in principle the opposite of electrolysis.

  • Inside an electrochemical cell or battery are chemicals that react together to produce electricity i.e. the cell produces a potential difference (p.d. or voltage).

  • The reactants constitute a supply of chemical potential energy to be converted into electrical energy.

  • A cell will produce a voltage until on of the reactants is used up.

  • Batteries consist of two or more cells connected in series to increase the p.d. voltage.

  • An oxidation-reduction (redox) reaction occurs to produce products and energy is given out because it is an exothermic reaction, BUT the energy is released as electrical energy NOT heat energy so the system shouldn't heat up.

  • A electrochemical simple cell can be made by dipping two different pieces of metal (of different reactivity) into a solution of ions e.g. a salt or dilute acid which will act as an electrolyte.

    • The arrangement is shown in the simple diagram of simple cell (above right), where the two pieces of metal are connected via a voltmeter to complete the electrical circuit.

    • If you connect several cells together in series, the voltage is increased (cell 1 V + cell 2 V etc. = total p.d.) and you might light up a bulb! having made a crude battery!

    • If you use the same metal for both strips, they 'cancel' each other out, so no potential difference (voltage) so no current of electrical energy.

  • All you need is a solution of charged positive and negative particles called ions e.g. sodium Na+, chloride Cl, hydrogen H+, sulphate SO42– etc.

  • The greater the difference in reactivity, the bigger the voltage produced. However this is not a satisfactory 'battery' for producing even a small continuous current.

  • BUT a simple demonstration cell can be made by dipping strips of magnesium and copper into an a salt solution (or a dilute acid) and connecting them via a voltmeter (e.g. as in diagram) and a voltage is readily recorded.

    • The electrode half-reactions are:

      • at the (+) electrode 2H+(aq) + 2e ==> H2(g) (hydrogen ions reduced)

        • the hydrogen ion - hydrogen half-equation

        • here the copper is inert and the hydrogen ions come from water.

      • at the () electrode Mg(s) – 2e ==> Mg2+(aq) (magnesium atoms oxidised)

        • the magnesium atom  - magnesium ion half-equation

        • the magnesium dissolves into solution by a chemical reaction

      • Each of the above equations is called a half–cell reaction, because that's what it is – half the chemical change.

      • So, overall the overall redox reaction is ...

        • 2H+(aq) + Mg(s) ==> Mg2+(aq) + H2(g) 

      • and the electrons from the oxidation of the magnesium move round through the magnesium strip, along the external wire to the copper electrode. In this case the copper strip just acts as an electrical connection and doesn't chemically change, but hydrogen ions from the water (or acid).

      • Note the (+) and (–) polarity of the electrodes in a cell, is the opposite of electrolysis because the process is operating in the opposite direction i.e.

        • in electrolysis electrical energy induces chemical changes,

        • but in a cell, chemical changes produce electricity.

  • One of the first practical batteries is called the 'Daniel cell' which is illustrated below.

    • The Daniel Cell

    • This 'voltaic 'or galvanic' electrochemical cell uses a half–cell of copper dipped in copper(II) sulphate,

    • and in electrical contact with another half–cell of zinc dipped in zinc sulphate solution.

    • The zinc is the more reactive, and is the negative electrode, releasing electrons because

      • on it zinc atoms lose electrons to form zinc ions, Zn(s) ==> Zn2+(aq) + 2e

        • the zinc atom - zinc ion half-equation

    • The less reactive metal copper, is the positive electrode, and gains electrons from the negative electrode through the external wire connection and here ..

      • the copper(II) ions are reduced to copper atoms, Cu2+(aq) + 2e ==> Cu(s)

        • the copper ion - copper atom half-equation

    • Overall the reactions is: Zn(s) + CuSO4(aq) ==> ZnSO4(aq) + Cu(s)

      • or ionically: Zn(s) + Cu2+(aq) ==> Zn2+(aq) + Cu(s)

    • The overall reaction is therefore the same as displacement reaction, and it is a redox reaction involving electron transfer and the movement of the electrons through the external wire to the bulb or voltmeter etc. forms the working electric current.

  • A simple voltaic cell from two metal strips dipped in an acid or salt solution (c) doc bThe cell voltage can be predicted by subtracting the less positive voltage from the more positive voltage:

    • e.g. referring to the list of electrode potentials on the right and choosing two different metals coupled together in the electrolyte solution ...

    • a magnesium and copper cell will produce a voltage of (+0.34) – (–2.35) = 2.69 Volts

    • or an iron and tin cell will only produce a voltage of (–0.15) – (–0.45) = 0.30 Volts.

    • Note (i) the bigger the difference in reactivity, the bigger the cell voltage produced

    • and (ii) the 'half–cell' voltages quoted in the diagram are measured against the H+(aq)/H2(g) system which is given the standard potential of zero volts.

  • Cells or batteries are useful and convenient portable sources of energy for torches, radios, shavers and other gadgets BUT they are expensive compared to what you pay for 'mains' electricity. On the other hand you have no choice for a car battery!

  • With rechargeable cells and batteries, it is possible to input electrical energy (via a charger) and reverse the chemistry that produced the electricity in the first place. The energy is then stored again as chemical potential energy and the battery can used again.

  • In non-rechargeable cells and batteries the chemical reactions must stop when one of the reactants has been used up.

    • You can't produce electricity if one of the reactants is no longer present!

    • The common zinc-carbon and acid paste battery comes into this category, so don't try and recharge it!

    • AND most alkaline batteries are non-rechargeable too.

  • See also 11. Fuel Cells e.g. the hydrogen - oxygen fuel cell

  • Electrolysis and cell theory for Advanced Level Chemistry Students

 

ELECTROCHEMISTRY INDEX:  1. INTRODUCTION to electrolysis - electrolytes, non-electrolytes, electrode equations, apparatus 2. Electrolysis of acidified water (dilute sulfuric acid) and some sulfate salts and alkalis 3. Electrolysis of sodium chloride solution (brine) and bromides and iodides 4. Electrolysis of copper(II) sulfate solution and electroplating with other metals e.g. silver 5. Electrolysis of molten lead(II) bromide (and other molten ionic compounds) 6. Electrolysis of copper(II) chloride solution 7. Electrolysis of hydrochloric acid 8. Summary of electrode equations and products 9. Summary of electrolysis products from various electrolytes 10. Simple cells (batteries) 11. Fuel Cells e.g. the hydrogen - oxygen fuel cell 12. The electrolysis of molten aluminium oxide - extraction of aluminium from bauxite ore & anodising aluminium to thicken and strengthen the protective oxide layer 13. The extraction of sodium from molten sodium chloride using the 'Down's Cell' 14. The purification of copper by electrolysis 15. The purification of zinc by electrolysis 16. Electroplating coating conducting surfaces with a metal layer 17. Electrolysis of brine (NaCl) for the production of chlorine, hydrogen & sodium hydroxide AND 18. Electrolysis calculations


Electrolysis Quiz (GCSE 9-1 HT Level (harder)

Electrolysis Quiz (GCSE 9-1 FT Level (easier)

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