10. Simple Cells and batteries

(A) Introduction

How a simple cell can be used as a battery is explained, using the different reactivities of two metal strips.

How can you make a simple battery, how can you use a simple cell to investigate the reactivity series of metals.

What is the difference between rechargeable and non-rechargeable batteries, how to make a simple copper-zinc cell, how to make a simple copper-magnesium cell.

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 and 5-1) science courses.

• 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, NOT thermal energy, so the system shouldn't heat up.

• A 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 - can be a salt solution of dilute acid.

  • 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.

  • If the metals have different reactivities, then an electrical current is generated as long as the circuit is complete as illustrated above on the right.

• You need is a solution of charged positive and negative particles called ions e.g. sodium Na⁺, chloride Cl^–, hydrogen H⁺, sulphate SO₄^2– 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!

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

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

  • The electrolyte here is a non-reactive aqueous of a salt, but it will work with a very dilute sulfuric acid solution.

  • (You can experiment with different electrolytes and see which one gives the greatest p.d. in volts,)

  • The electrode half-reactions are:

    • at the (+) electrode 2H⁺_(aq) +  2e^–  ===> H_2(g)

      • (hydrogen ions reduced on the surface of the copper because there are no copper ions to be reduced)

      • above is the hydrogen ion - hydrogen half-equation

      • here the copper is inert and the hydrogen ions come from water/acid and end up as bubbles of hydrogen.

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

    • at the (–) electrode Mg_(s) – 2e^– ===>  Mg²⁺_(aq)

      • (magnesium atoms oxidised, the electrons run round the wire to reduce the hydrogen ions from water or acid)

      • above is the magnesium atom  - magnesium ion half-equation

      • the magnesium dissolves into solution by the electrode 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) + H_2(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.

    • The electrode potential of each metal is a measure of its chemical reactivity - the more negative or less positive, the more reactive the metal and the greater the difference in the two metal strips, the larger the p.d. in volts created.

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

      • See section (D) on how to predict the voltage (p.d.) 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 - see end of section (D).

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(C) An early practical battery cell

• As we have seen above, the simplest cell to generate electricity can be made by dipping two externally connected pieces of different metals into an electrolyte solution of a non-reactive salt.

•  e.g. connecting strips of zinc and copper (plus voltmeter) and placing in the electrolyte of zinc sulfate, .

• 

  • This simple cell system 'sort of' works for a few minutes and then the voltage drops away..

  • So this set-up is of not practical enough for any use!

  • It won't work effectively as a battery with just one electrolyte.

  • SO, you do need something a bit more sophisticated than this simple cell and the actual set-up in principle for the Daniel Cell is shown below, one of the first simple, but effective, batteries used in laboratories as an electrical supply for other experiments.

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

• 

  • The 'fuel' is effectively zinc metal and copper(II) sulfate solution which get consumed when the battery is working to generate a constant stream of d.c. electrical current.

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

  • and in electrical contact with another half–cell of zinc metal 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 acts as an inert electrode.

    • Instead copper(II) ions gain electrons from the negative electrode through the external wire connection and are reduced to copper metal ..

    • the copper ions are reduced to copper atoms: Cu²⁺_(aq) + 2e^– ===> Cu_(s)

      • the copper ion - copper atom half-equation

  • Overall the reactions is:

    • Zn_(s)  +  CuSO_4(aq)  ===>  ZnSO_4(aq)  +  Cu_(s)

    • or ionically: Zn_(s)  +  Cu²⁺_(aq)  ===>  Zn²⁺_(aq)  +  Cu_(s)

    • It is an exothermic reaction, BUT here, there is no temperature rise, because the energy is released as electrical energy carried by the flow of electrons.

    • The theoretical p.d. created is 0.34 - (-0.76) = 1.10 V.

    • The electrode potential of each metal is a measure of its chemical reactivity - the more negative or less positive, the more reactive the metal.

  • 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.

    • You can make a similar cell system using magnesium/magnesium sulfate instead of zinc/zinc sulfate which would give a voltage of 0.34 - (-2.35) = 2.69 V.

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

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(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 subtracting most negative from the least negative):

  • Technically, in more advanced chemistry, these individual electrode voltages are called the half-cell potentials.

  • So, by 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 V 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 hydrogen ion - hydrogen gas potential H⁺_(aq)/H_2(g) system which is given the arbitrary standard potential of zero volts (hydrogen/hydrogen ion potential = 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 copper - zinc cell can be simply set up with strips of the two metals dipped into a salt or dilute acid solution.

  • The predicted voltage would be copper potential - zinc potential

  • V = +0.34 - (-0.76) = 1.10 V

  • Note in practice that the p.d. (V) you measure for the initial minute might vary from electrolyte to electrolyte because the chemistry is not quite as simple as the diagram suggests.

  • Also, without the presence of copper ions you are more likely to measure a voltage of ~0.76 V because the copper may act as an inert electrode and hydrogen forming on its surface (0 - (-0.76) = 0.76 V.

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

  •  

  • A copper - magnesium cell can be simply set up with strips of the two metals dipped into non-reactive salt or dilute acid solution.

  • The predicted voltage would be copper potential - magnesium potential

  • V = +0.34 - (-2.35) = 2.69 V

  • Again, without the presence of copper ions you are more likely to measure a voltage of ~2.35 V b because the copper may act as an inert electrode and hydrogen forming on its surface (0 - (-2.35) = 2.35 V.

  • A case of setting up two cells in series to produce a bigger voltage

  • 

  • The predicted cell voltages are 2.35 V (copper not involve chemically) and 1.10 V for the two cells.

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

  • Practical re-chargeable batteries, like a car battery, are made up several 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.

  • Before modern electricity supplies were available, lots of Daniel Cells were linked together in series to produce much higher voltages and were known as 'voltaic piles' see https://en.wikipedia.org/wiki/Voltaic_pile

  • Extra note on the reactivity series of metals

    • A simple cell of a carbon rod and a metal strip.

    • This is a simple experiment to investigate the reactivity series of metals - the more reactive the metal, the greater the p.d. in volts, but make sure you always connect the voltmeter the same way round and produce a positive value for e.g. magnesium.

    • A carbon rod (graphite) can be used as a bench mark chemically inert electrode and the voltages measured for a series of other metals paired with it to get a partial metal reactivity series.

    • This is the simplest experiment I know to use a simple cell to determine a metal reactivity series.
    • It should work well enough with very dilute sulfuric acid, with other metals like tin or nickel, readings can be close and trend not clear, but it might show Al in its correct place despite the inhibiting oxide layer.

    • You should obtain a series of increasing p.d values (V) e.g. Mg > Zn > Fe >Cu, effectively a measure of the reactivity of the metal.

    • but the actual values may depend on the electrolyte used - you can try other electrolytes like aqueous sodium sulfate solution.

(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!

    • In this type of battery, the zinc reacts with an acid paste and the hydrogen formed is oxidised to water with an oxidising agent.

  • AND most alkaline batteries are non-rechargeable too.

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

• Electrolysis and cell-battery theory-examples for Advanced Level Chemistry Students