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GCSE Chemistry notes: Describing and explaining how simple batteries work

(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. The chemistry of simple cells & batteries

Sub-index for this page

(A) Introduction to cells and batteries

(B) A simple cell experiment to investigate the effects of using different pairs of metals

(C) An early practical battery cell

(D) More on investigation experiments and how to predict the cell voltage

(E) Practical batteries for commercial and domestic use - chargeable and non-rechargeable

See also 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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10. Simple Cells and batteries


(A) Introduction

  • 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) and an electrical current flows - electrons in the wire and ions in the solution - the current produced is d.c., it only flows in one direction.

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

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

  • An oxidation-reduction (redox) reactions occurs at electrodes to produce products and energy is given out because it is an exothermic reaction,

    • BUT the energy is released as electrical energy thermal energy so the system shouldn't heat up.

  • A simple voltaic cell from two metal strips dipped in an acid or salt solution (c) doc bA simple electrochemical cell can be made by dipping two different pieces of metal (must be of different reactivity - different potential), connected by a wire, into a solution of ions e.g. a salt or dilute acid which will act as an electrolyte.

    • The electrolyte is a solution of charged particles - ions, that can carry an electric current.

    • The external wire and voltmeter completes the circuit - as in physics!

    • The two pieces of metal can be held in crocodile clips and acts as electrodes - the electrical contacts with the electrolyte solution - at least one must react, one may be inert, but they both usually react as part of the electrochemical cell chemistry.

    • The arrangement is shown in the simple diagram of simple cell (right)

    • If you connect several cells together in series, the voltage is increased.

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

  • The greater the difference in reactivity of the two metals, the bigger the cell voltage produced.

    • If you use the same metal for both strips, their chemical potentials 'cancel' each other out, so no potential difference (voltage = 0 V) so no current of electrical energy.

    • If you connect several cells together, identical or different, you can add up the individual cell voltages to give the total p.d. in volts  AND you might light up a bulb! having made a crude battery!

 


(B) A simple cell experiment to investigate the effects of using different pairs of metals

  • a simple magnesium hydrogen electrochemical cellA simple demonstration cell can be made by e.g. dipping strips of magnesium and copper metals 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 electrolyte here is aqueous sodium chloride solution.

    • The electrode half-reactions are:

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

        • the hydrogen ion - hydrogen half-equation

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

        • You won't see a copper deposit on the magnesium.

      • 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) do.

      • 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.

      • Theoretically you can generate a p.d. of 2.35 V, unlikely, but it should work and give you a voltage!

        • See section (D) on how to predict the voltage the cell generates.

      • I think this might work with just a carbon rod instead of copper - check that out!

      • If so, you could establish a reactivity series of metals based on voltages produced keeping the carbon rod electrode constant and varying the metal electrode.


(C) An early practical battery cell

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

  • a simple zinc copper electrochemical cell diagram explained The simple cell system that works, but not practical for any use!

    • BUT, you do need something a bit more sophisticated than this simple cell.

    • The diagram below explains the chemistry behind one of the first practical battery systems.

  • 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.

    • In a working Daniel cell the two solutions are separated by a porous barrier that ions can diffuse through to complete the electrical circuit.

 

 


(D) More on investigation experiments and how to predict the cell voltage

  • The positive cell voltage can be predicted by subtracting the less positive voltage from the more positive voltage (or the most negative from the least negative):

    • A simple voltaic cell from two metal strips dipped in an acid or salt solution (c) doc be.g. referring to the list of electrode potentials of voltages 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 if the electrolyte used is copper sulfate solution.

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

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

    • (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 (hydrogen = 0.00 V).

    • (iii) If you swap the metal electrodes around, you reverse the sign of the cell voltage (p.d.) and the current flows in the opposite direction - your digital meter reading might change from + 0.30 V to -).30 V.

    • You should appreciate the electrode potential is a measure of the chemical potential energy of that metal to react by losing electrons - remember the theory behind the reactivity series of metals!

      • Therefore this series of electrode potentials is identical to the reactivity series of metals.

      • Hydrogen is given a standard arbitrary value of 0.00 V

    • a simple zinc copper electrochemical cell diagram explained

    • A copper - zinc cell can be simply set up with strips of the two metals dipped into copper sulfate solution.

    • The predicted voltage would be copper potential - zinc potential

    • V = +0.34 - (-0.76) = 1.10 V

    • The chemistry for this electrochemical cell is described in section (C) above.

    • a simple copper magnesium electrochemical cell diagram explained 

    • A copper - magnesium cell can be simply set up with strips of the two metals dipped into copper sulfate solution.

    • The predicted voltage would be copper potential - magnesium potential

    • V = +0.34 - (-2.35) = 2.69 V

    • The magnesium gets oxidised (electron loss) and the copper ions get reduced (electron gain)

      • Half equations

        • Mg(s) – 2e ==> Mg2+(aq)

        • Cu2+(aq) + 2e ==> Cu(s)

        • the overall ionic equation for the cell reaction is: Mg(s) + Cu2+(aq) ==> Mg2+(aq) + Cu(s)

    • A case of setting up two cells in series

    • making a battery from connecting simple cells in series

    • The predicted cell voltages are 2.35 (copper not involve chemically) and 1.10.

    • Theoretically this more complex system should generate a total p.d of 2.35 + 1.10 = 3.45 V

    • Practical batteries, like a car battery, are made up of two or more cells connected in series to increase the working voltage.

    • In the lab, a class could put several similar simple cells together, wired in series, and see what higher voltages you could generate.


(E) Practical batteries for commercial and domestic use - rechargeable and non-rechargeable

  • The simple cells described above do not make a satisfactory 'battery' for producing even a small continuous d.c. current.

    • So the batteries you buy in shops are a bit more complicated.

  • 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!

    • Its all changed to the 'product' and there is no longer any chemical potential energy to be transferred as useful work - electrical energy.

    • 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

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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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