Chapter 6: Electrolysis
In the autumn of 1807, a young chemist named Humphry Davy was working in the basement laboratory of the Royal Institution in London, in front of the largest battery anyone had yet built — hundreds of metal plates stacked in acid and wired into one enormous pile.
Davy was after something nobody had ever seen. Chemists knew that potash, the white solid left behind when wood ash is washed and the water boiled away, held an unknown metal locked inside it. No one had managed to prise it out. Heating did not work. No other substance was greedy enough to take it.
Davy first dissolved potash in water and passed his current through the solution. He got bubbles of gas, and nothing more. So he tried again. This time he melted the potash and dipped his wires straight into the hot liquid. Almost at once, small silvery globules appeared at the wire joined to the negative end of the battery. They shone like mercury for a moment, then caught fire and burned with a bright lilac flame.
He had made potassium. Within days, the same method gave him sodium out of soda. Electricity, it turned out, could take apart compounds that no chemical reagent could touch. What follows in this chapter is how it does that.
Diagram
Fig. 6.0-1 — Davy’s Experiment On Molten Potash Simple period-style illustration: a tall voltaic pile of alternating metal discs on the left, two wires running from its two ends to a small dish of hot molten potash on a stand; show tiny bright globules of metal collecting on the wire from the negative end, and a bubble or two at the other wire; label the pile (battery), the wire from the negative end, the wire from the positive end, molten potash, and globules of potassium.
6.1 Conductors, Non-Conductors And The Passage Of Current
Davy’s two wires did nothing on their own. Everything depended on what he dipped them into. Water gave him gas, molten potash gave him metal, and a dry lump of the same potash would have given him nothing at all.
That is where this chapter begins. Some substances let an electric current pass through them and some do not; and of those that do let it pass, some are changed by it while others come out exactly as they went in. Sorting substances into these groups is the first step.
6.1.1 Conduction In Metals And Non-Metals
Most metallic elements — copper, aluminium, iron — and all alloys allow an electric current to pass through them. They do this whether they are solid or molten, and they come through the experience unchanged. A copper wire that has carried current for a year is still copper. Substances that behave in this way are called conductors of electricity.
Non-metals behave differently. Sulphur and phosphorus do not conduct at all. The one striking exception is graphite, which conducts freely. Non-metals, and materials such as rubber, glass and dry wood, are called non-conductors or insulators.
Why should metals be so different? In a metal, the outermost electrons of each atom are held only loosely, and they drift freely through the whole piece of metal. Conduction in a conductor is nothing more than the flow of these mobile electrons. An insulator has no such free electrons, so there is nothing inside it that can move when a battery is connected, and no current flows.
You can sort substances yourself with the simple circuit shown in Fig. 6.1-1.
Diagram
Fig. 6.1-1 — A Circuit For Testing Whether A Substance Conducts Simple test circuit drawn as a loop: a battery of two cells, a switch, a small bulb, and two carbon rods held apart in a beaker; label battery, switch, bulb, carbon rods (terminals) and the substance under test in the beaker. Show a second small inset of the same circuit with the two carbon rods pressed against a strip of metal instead of dipping into a beaker.
6.1.2 Electrolytes And Non-Electrolytes
Chemical compounds behave in a third way, and it is this behaviour that the rest of the chapter is about. A compound such as copper chloride will not conduct at all as a dry solid. Dissolve it in water, or melt it, and it conducts — but it does not survive the experience. It is taken apart.
Pass a current through a solution of copper chloride and you can watch it happen. A layer of reddish brown copper builds up on one terminal, and a sharp, choking greenish yellow gas comes off the other. The compound has been broken into its two elements.
copper chloride solution → copper + chlorine
\[\mathrm{CuCl_2}\text{(aq)} \xrightarrow{\ \text{electric current}\ } \mathrm{Cu}\text{(s)} + \mathrm{Cl_2}\text{(g)}\]
Compounds that behave in this way are called electrolytes.
Other compounds do nothing at all. Sugar dissolves readily in water, but the solution will not light the bulb, whether the sugar is solid or dissolved. Nothing is deposited and nothing is given off. Such compounds are called non-electrolytes.
6.1.3 Metallic Conduction Compared With Electrolytic Conduction
The two kinds of conduction differ in what carries the charge, in what states the substance must be in, and in whether anything new is made. Table 6.1 sets them side by side.
The last row is the one to hold on to. Everything that follows in this chapter — metals deposited, gases given off, articles plated with silver — comes out of that single difference.
6.2 Strong Electrolytes, Weak Electrolytes And Non-Electrolytes
Electrolytes are not a small or exotic group. Between them they cover the three great families of compounds you already know — the acids, the bases and the salts. Molten salts, solutions of salts in water, and solutions of acids and alkalis are all electrolytes.
Some Common Electrolytes
- Acids — \(\mathrm{H_2SO_4}\), \(\mathrm{HNO_3}\), \(\mathrm{H_3PO_4}\), \(\mathrm{HCl}\)
- Bases — \(\mathrm{NaOH}\), \(\mathrm{KOH}\), \(\mathrm{NH_4OH}\), \(\mathrm{Ca(OH)_2}\)
- Salts — \(\mathrm{NaCl}\), \(\mathrm{CuSO_4}\), \(\mathrm{PbBr_2}\)
They do not all conduct equally well, though, and that difference divides them into two types.
6.2.1 Strong And Weak Electrolytes
Put the circuit of Fig. 6.1-1 into two beakers in turn — dilute hydrochloric acid in one, acetic acid of the same strength in the other. The bulb glows brightly in the first and only dimly in the second. Both liquids conduct, but one carries far more current than the other.
The reason lies in how completely the compound has broken up into free ions. A strong electrolyte is one that is almost completely split into ions in the fused or aqueous state, so its solution contains, in effect, only free mobile ions. A weak electrolyte is only partly split, so its solution contains unsplit molecules alongside the ions. Fewer ions means fewer charge carriers, and a dimmer bulb.
The extent of this splitting is measured by the degree of dissociation, given the symbol \(\alpha\).
A strong electrolyte has a high degree of dissociation; a weak electrolyte has a low one. Table 6.2 gathers the differences.
6.2.2 Non-Electrolytes And Their Examples
A non-electrolyte has no ions in it at all, in solution or out of it. It is made of molecules and stays made of molecules, so there is nothing charged that can move towards a terminal and nothing that can be discharged there.
Common non-electrolytes include distilled water, alcohol, kerosene, carbon disulphide, cane sugar, benzene, glucose and urea.
Distilled water on that list is worth a second look. Water does split into ions to an extremely small extent, but the number is so tiny that pure water conducts almost not at all, and it is treated as a non-electrolyte. That faint ionisation becomes important later, in Section 6.8.
Molecules Only, Ions Only, Or Both
Sorting the three groups by what their solutions actually contain is a useful check, and it comes straight out of what you have just read.
- Only molecules — a non-electrolyte such as cane sugar solution or carbon tetrachloride.
- Only ions (very nearly) — a strong electrolyte such as sodium chloride solution or molten lead bromide.
- Ions as well as molecules — a weak electrolyte such as acetic acid or ammonium hydroxide solution.
6.2.3 Acids, Bases And Salts As Electrolytes
Acids and alkalis are covalent or partly covalent substances, and they produce their ions only on being dissolved in water. Bases and salts are built of ions already, and they set those ions free either on melting or on dissolving. Either way the result is the same — a liquid full of free mobile ions, ready to carry current.
Whether the compound counts as strong or weak then depends on its degree of dissociation. Table 6.3 gives the equation for each, with the strong ones grouped separately from the weak.
Table 6.3 — Ionisation And Dissociation Of Some Acids, Bases And Salts
| Acids | Bases | Salts | |
|---|---|---|---|
| Strong electrolytes | Hydrochloric acid \(\mathrm{HCl} \rightarrow \mathrm{H^{+}} + \mathrm{Cl^{-}}\) [aq.] Nitric acid \(\mathrm{HNO_3} \rightarrow \mathrm{H^{+}} + \mathrm{NO_3^{-}}\) [aq.] Sulphuric acid \(\mathrm{H_2SO_4} \rightarrow 2\,\mathrm{H^{+}} + \mathrm{SO_4^{2-}}\) [aq.] |
Potassium hydroxide \(\mathrm{KOH} \rightarrow \mathrm{K^{+}} + \mathrm{OH^{-}}\) [aq.] Sodium hydroxide \(\mathrm{NaOH} \rightarrow \mathrm{Na^{+}} + \mathrm{OH^{-}}\) [aq.] Lithium hydroxide \(\mathrm{LiOH} \rightarrow \mathrm{Li^{+}} + \mathrm{OH^{-}}\) [aq.] |
Lead bromide \(\mathrm{PbBr_2} \rightarrow \mathrm{Pb^{2+}} + 2\,\mathrm{Br^{-}}\) [molten] Copper chloride \(\mathrm{CuCl_2} \rightarrow \mathrm{Cu^{2+}} + 2\,\mathrm{Cl^{-}}\) [aq.] Silver nitrate \(\mathrm{AgNO_3} \rightarrow \mathrm{Ag^{+}} + \mathrm{NO_3^{-}}\) [aq.] |
| Weak electrolytes | Acetic acid \(\mathrm{CH_3COOH} \rightleftharpoons \mathrm{CH_3COO^{-}} + \mathrm{H^{+}}\) [aq.] Formic acid \(\mathrm{HCOOH} \rightleftharpoons \mathrm{HCOO^{-}} + \mathrm{H^{+}}\) [aq.] |
Ammonium hydroxide \(\mathrm{NH_4OH} \rightleftharpoons \mathrm{NH_4^{+}} + \mathrm{OH^{-}}\) [aq.] Calcium hydroxide \(\mathrm{Ca(OH)_2} \rightleftharpoons \mathrm{Ca^{2+}} + 2\,\mathrm{OH^{-}}\) [aq.] |
Ammonium carbonate \(\mathrm{(NH_4)_2CO_3} \rightleftharpoons 2\,\mathrm{NH_4^{+}} + \mathrm{CO_3^{2-}}\) [aq.] Lead acetate \(\mathrm{(CH_3COO)_2Pb} \rightleftharpoons 2\,\mathrm{CH_3COO^{-}} + \mathrm{Pb^{2+}}\) [aq.] |
Notice the two kinds of arrow. A single arrow says the change goes essentially to completion, which is what happens with a strong electrolyte. The double arrow of a weak electrolyte says that ions and molecules exist side by side, with molecules splitting and ions rejoining all the time.
Salts as a family are strong electrolytes, and they dissociate completely to give solutions that may be acidic, basic or neutral. A few are exceptions — salts that are more volatile than the rest, or that do not dissociate completely, behave as weak electrolytes. Mercury(II) chloride, \(\mathrm{HgCl_2}\), and ammonium carbonate are the usual examples.
6.3 The Language Of Electrolysis
Everything that follows in this chapter is described with about a dozen words — cell, electrode, anode, cathode, ion, cation, anion, oxidation, reduction. Each of them names something real that you could point at in a working cell. This section fixes what each one means, so that later sections can be read without stopping.
6.3.1 Electrolysis And The Electrolytic Cell
When an electrolyte in its molten or aqueous state has a direct current passed through it, the compound is torn apart into the ions it is built from, and those ions are turned back into ordinary atoms or molecules at the two terminals. That whole process is electrolysis.
The apparatus in which this is done is the electrolytic cell, also called a voltameter. It is a vessel of some non-conducting material holding the electrolyte, with two conducting plates or rods dipping into it and joined to a battery. The battery is essential: nothing happens on its own, and the cell has to be driven from outside by a direct-current supply. What such a cell does, in energy terms, is convert electrical energy into chemical energy. A simple cell is shown in Fig. 6.3-1.
Diagram
Fig. 6.3-1 — A Simple Electrolytic Cell (Voltameter) Rectangular glass vessel about two-thirds full of electrolyte, with two flat rectangular electrodes hanging vertically into it from above and held by a lid; connecting wires run from the tops of the two electrodes to the two terminals of a battery drawn above, with the positive and negative terminals marked. Label battery, positive terminal, negative terminal, anode, cathode, electrolytic cell and electrolyte.
6.3.2 The Electrochemical Cell
An electrochemical cell does the opposite job. It uses a chemical reaction to produce an electric current, converting chemical energy into electrical energy, and it needs no outside supply because it is itself the supply. The simple voltaic cell and the Daniell cell are examples.
Both kinds of cell contain two electrodes in an electrolyte, which is why they are easily confused. What separates them is the direction in which the energy travels.
6.3.3 Electrodes: Anode And Cathode
The two plates or rods dipping into the electrolyte are the electrodes. They may be metal plates, metal wires, graphite rods or gas carbon rods, and the current enters and leaves the electrolytic cell through them.
They are told apart by which terminal of the battery they are joined to. The anode is the electrode connected to the positive terminal, and the cathode is the electrode connected to the negative terminal.
It helps to follow the electrons. Once the current is turned on, the battery pushes electrons round the external wire from the anode to the cathode. So the cathode becomes the negative electrode, the one by which electrons enter the electrolyte, and the anode becomes the positive electrode, the one by which electrons leave the electrolyte. Everything that happens chemically in the cell follows from those two facts.
Table 6.4 sets the two electrodes against each other. The last row uses two words that are defined a little further on, in Section 6.3.5.
6.3.4 Ions: Cations And Anions
An ion is an atom, or a group of atoms, that carries a positive or a negative charge. The charge it carries is equal to the valency of that atom or group — a sodium atom of valency one gives \(\mathrm{Na^{+}}\), an aluminium atom of valency three gives \(\mathrm{Al^{3+}}\), and the sulphate group of valency two gives \(\mathrm{SO_4^{2-}}\).
Ions of the two signs are given separate names by where they travel during electrolysis. A cation is a positively charged ion; opposite charges attract, so it is pulled towards the negative cathode. An anion is a negatively charged ion, pulled towards the positive anode.
6.3.5 Oxidation And Reduction As Loss And Gain Of Electrons
You have met oxidation as the addition of oxygen or removal of hydrogen, and reduction as the removal of oxygen or addition of hydrogen. Both descriptions still hold, and examples of each follow below. But neither of them can describe what happens at an electrode, where there is often no oxygen and no hydrogen anywhere in sight. What there always is, is a transfer of electrons — and that is what the electronic definitions are built on.
\[\mathrm{Zn} - 2e^{-} \rightarrow \mathrm{Zn^{2+}}\] \[\mathrm{Na} - e^{-} \rightarrow \mathrm{Na^{+}}\] \[\mathrm{Fe^{2+}} - e^{-} \rightarrow \mathrm{Fe^{3+}}\] \[\mathrm{S^{2-}} - 2e^{-} \rightarrow \mathrm{S}\]
Read the third of these carefully. Nothing is added and nothing is removed; an ion simply becomes more positive by shedding one electron. That is oxidation all the same.
\[\mathrm{Cu^{2+}} + 2e^{-} \rightarrow \mathrm{Cu}\] \[\mathrm{Fe^{3+}} + e^{-} \rightarrow \mathrm{Fe^{2+}}\] \[\mathrm{Sn^{2+}} + 2e^{-} \rightarrow \mathrm{Sn}\] \[\mathrm{S} + 2e^{-} \rightarrow \mathrm{S^{2-}}\]
The Older Descriptions, Still In Use
Oxidation is also defined as a chemical process involving the addition of oxygen or the removal of hydrogen.
carbon + oxygen → carbon dioxide
\[\mathrm{C} + \mathrm{O_2} \rightarrow \mathrm{CO_2}\]
\[2\,\mathrm{Mg} + \mathrm{O_2} \rightarrow 2\,\mathrm{MgO}\]
Here carbon and magnesium are oxidised by the gain of oxygen. In the next reaction, hydrogen sulphide is oxidised to sulphur by the loss of hydrogen.
\[\mathrm{H_2S} + \mathrm{Cl_2} \rightarrow 2\,\mathrm{HCl} + \mathrm{S}\]
Reduction, in the same older language, is the removal of oxygen or the addition of hydrogen.
\[\mathrm{CuO} + \mathrm{H_2} \rightarrow \mathrm{Cu} + \mathrm{H_2O}\] \[\mathrm{ZnO} + \mathrm{C} \rightarrow \mathrm{Zn} + \mathrm{CO}\]
Copper(II) oxide and zinc oxide are reduced to copper and zinc by the loss of oxygen. In the two reactions below, chlorine is reduced to hydrogen chloride by the gain of hydrogen.
\[\mathrm{Cl_2} + \mathrm{H_2S} \rightarrow \mathrm{S} + 2\,\mathrm{HCl}\] \[3\,\mathrm{Cl_2} + 2\,\mathrm{NH_3} \rightarrow \mathrm{N_2} + 6\,\mathrm{HCl}\]
6.3.6 Oxidising Agents And Reducing Agents
If one substance loses electrons, something must be taking them. An oxidising agent is a substance that oxidises another substance, either by accepting electrons from it, or by providing it with oxygen or an electronegative ion, or by removing hydrogen or an electropositive ion from it. In the reactions above, oxygen and chlorine are the oxidising agents.
Some Oxidising Agents
- Solids — manganese dioxide, red lead, lead dioxide
- Liquids — hydrogen peroxide, concentrated nitric acid, concentrated sulphuric acid, bromine
- Gases — oxygen, ozone, chlorine, sulphur dioxide
A reducing agent works the other way. It reduces another substance by providing electrons, or by providing hydrogen or an electropositive ion, or by removing oxygen or an electronegative ion from it. In the reactions above, hydrogen, carbon, hydrogen sulphide and ammonia are the reducing agents.
Some Reducing Agents
- Solids — carbon, metals such as zinc, aluminium, copper and sodium, stannous chloride, glucose
- Liquids — hydrogen peroxide, hydrogen iodide, hydrogen bromide
- Gases — hydrogen sulphide, carbon monoxide, sulphur dioxide
Two names appear on both lists, and that is not a slip. Hydrogen peroxide and sulphur dioxide can each behave either way, depending on what they are put with. Which role a substance plays is decided by the reaction it finds itself in, not by the substance alone.
6.4 How Compounds Come To Contain Free Ions
A cell can only work if the liquid in it contains ions that are free to move. Table 6.1 said as much, and Table 6.2 turned on how many such ions there are. The question this section answers is where those free ions come from, and why the same compound supplies them in one state and not in another.
6.4.1 Arrhenius And The Theory Of Electrolytic Dissociation
In 1887 the Swedish chemist Svante Arrhenius put forward an ionic theory to account for the conducting power of solutions. Its main points are these.
- An electrolyte, on dissolving in water, breaks up into free mobile ions — cations carrying positive charge and anions carrying negative charge — and this allows an electric current to flow through it.
- Every one of these ions carries an electric charge, and it is their movement that carries the current through the solution.
- How well the electrolyte conducts depends on the concentration of ions in the solution.
- The number of positive charges in the solution equals the number of negative charges, so the solution as a whole is neutral and stands in electrolytic equilibrium. An equilibrium is also set up between the molecules that have not split and the ions they produce.
- In non-electrolytes such as sugar solution or benzene there is no ionisation at all, so only molecules are present in the solution.
- The extent to which an electrolyte breaks up into ions is its degree of dissociation, met in Section 6.2.1.
Where The Modern Picture Differs
Arrhenius pictured water as the agent that made the ions — the electrolyte went in as molecules and water broke them up. The modern view corrects this. An electrolyte such as sodium chloride is built of ions already, even as a dry solid; those ions are simply locked in place by strong electrostatic forces and cannot move. What water does is not to create the ions but to break the forces holding them, giving ions that were already there the freedom to move.
6.4.2 Conduction In Electrovalent Compounds
An electrovalent compound is an aggregate of oppositely charged ions, each held in its position by the strong electrostatic force of attraction of its neighbours. In the solid state those ions cannot travel anywhere. This is why a dry lump of sodium chloride will not conduct, even though it is packed with ions from end to end — the ions are present but immobile.
Melt the compound, or dissolve it in water, and the forces of attraction between the ions are broken. The ions can then move about freely, and the liquid conducts.
Not every electrolyte is electrovalent, though. Covalent compounds whose molecules are polar also form ions when dissolved in water, and their solutions conduct. Hydrogen chloride, ammonia and water itself are of this kind.
6.4.3 Electrolytic Dissociation Of Sodium Chloride
The splitting apart of a compound that is ionic to begin with has its own name — electrolytic dissociation.
Take sodium chloride through it step by step.
Solid sodium chloride is an electrovalent compound, and its crystal structure is a regular array of sodium ions, \(\mathrm{Na^{+}}\), and chloride ions, \(\mathrm{Cl^{-}}\). In the solid state these ions are not free; they are held to one another by electrostatic attraction, as Fig. 6.4-1 shows, and so the solid does not conduct electricity. The absence of free ions, not the absence of ions, is the reason.
Diagram
Fig. 6.4-1 — Ions Held In Place In Solid Sodium Chloride Two large circles side by side, one labelled Na⁺ and one labelled Cl⁻, drawn touching, with short dotted lines radiating outwards from each to indicate attraction to further ions beyond the edge of the drawing and a dense band of dotted lines between the two circles. Caption note for the illustrator: the dotted lines represent the electrostatic force of attraction; no gap should appear between the two ions.
When the solid is dissolved in water or melted, it conducts. The two routes work differently.
- In the aqueous state, the water molecule is polar: its oxygen atom carries a slight negative charge and its hydrogen atoms a slight positive charge. The slightly negative oxygen atoms pull on the positively charged sodium ions, and the slightly positive hydrogen atoms pull on the negatively charged chloride ions. Pulled from both sides, the ions come away from the crystal and are free in solution, where they conduct electricity.
- In the molten state, no water is involved at all. The high temperature needed to melt the solid weakens the bond between the particles, and the ions are set free, as in Fig. 6.4-2.
Diagram
Fig. 6.4-2 — Ions Set Free In Solution Or In The Molten State The same two circles labelled Na⁺ and Cl⁻, now drawn well apart with the dotted lines between them broken and shown parting; add a small arrow on each ion showing it moving away from the other, and the label ‘weakening of the bond’ between them.
6.4.4 Ionisation Of Polar Covalent Compounds
Some covalent molecules contain no ions whatever, and yet their aqueous solutions conduct. Something must be making ions that were not there before, and that process is called ionisation.
Hydrogen chloride is the standard case, and it goes in three stages.
- As a gas, or as a pure liquid, hydrogen chloride is unionised. It does not conduct an electric current.
- Its molecule is polar covalent — the shared electrons are held closer to chlorine, so the hydrogen atom carries a slight positive charge and the chlorine atom a slight negative charge, written \(\mathrm{H^{\delta+}}\!-\!\mathrm{Cl^{\delta-}}\).
- Added to water, which is itself polar, the slightly negative oxygen atom of a water molecule pulls hard on the slightly positive hydrogen of the hydrogen chloride molecule — hard enough to take it away as a hydrogen ion and leave the chloride ion behind.
\[\mathrm{HCl} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^{+}} + \mathrm{Cl^{-}}\]
The hydrogen ion does not stay on its own. It combines with a water molecule to form the hydronium ion, \(\mathrm{H_3O^{+}}\), and it is this ion, together with the chloride ion, that carries the current through the solution. Fig. 6.4-3 shows the transfer.
Diagram
Fig. 6.4-3 — Ionisation Of Hydrogen Chloride In Water Left side: an HCl molecule drawn as two joined circles labelled H with δ+ and Cl with δ−; beside it a bent water molecule with the O labelled δ− and its two H atoms labelled δ+. A curved arrow runs from the H of HCl to the O of water. Right side of the arrow: a hydronium ion drawn as an O with three H atoms and an overall + charge, labelled hydronium ion, and separately a circle labelled Cl⁻, labelled chloride ion.
6.4.5 Ionisation Compared With Dissociation
The two words are often used loosely as though they meant the same thing. They do not, and Table 6.6 separates them.
6.5 Characteristics Of Electrolysis
Whatever the electrolyte and whatever the electrodes, every electrolysis behaves in the same general way. Eight features are worth setting out together, because the worked examples in Section 6.8 use them again and again.
The order of events is always the one shown in FC 6.1.
Flowchart
FC 6.1 — What Happens During An Electrolysis Electrolyte melted or dissolved → free mobile ions produced → direct current switched on → cations travel to the cathode and anions travel to the anode → at the cathode each cation gains electron(s) and is reduced; at the anode each anion loses electron(s) and is oxidised → neutral atoms formed, which may combine into molecules → products collected at the two electrodes
The Eight Characteristics
- Passing electricity through an electrolyte makes the positive ions, the cations, migrate towards the cathode, and the negative ions, the anions, migrate towards the anode.
- The number of electrons gained by the anode equals the number of electrons given out by the cathode. It has to be so: electrons are not made or destroyed in the cell, and none pile up anywhere in the circuit.
- The products of electrolysis are formed at the anode and the cathode themselves, because the exchange of electrons happens only at the surface of the electrodes.
- Only hydrogen gas and metals are liberated at the cathode. This follows from the first characteristic — only cations reach the cathode, and cations are either metal ions or hydrogen ions. Because they carry positive charge, these are called electropositive elements.
- Only non-metals are liberated at the anode, for the matching reason: only anions reach it. These are called electronegative elements.
- The mass of a substance produced at an electrode during electrolysis is proportional to the quantity of electricity passed through the electrolyte.
- The process of electrolysis is a redox reaction. Reduction and oxidation are not two separate events here but the two halves of one; every electron handed to a cation at the cathode is an electron that some anion has surrendered at the anode.
- Alternating current brings about no chemical change in an electrolyte and will not drive an electrolysis. Electrolysis is caused by direct current only. With alternating current the two electrodes swap sign many times each second, so an ion is pulled first one way and then the other; whatever is deposited during one half of the cycle is dissolved back during the next, and nothing accumulates.
Electrolysis As A Redox Reaction
Take molten sodium chloride as the example. In the fused state it dissociates into its two ions.
\[\mathrm{NaCl} \rightarrow \mathrm{Na^{+}} + \mathrm{Cl^{-}}\]
Sodium ions travel to the cathode. Each takes one electron and becomes a neutral sodium atom, which is reduction.
\[\text{At the cathode: } \mathrm{Na^{+}} + e^{-} \rightarrow \mathrm{Na} \quad \text{(reduction)}\]
Chloride ions travel to the anode. Each gives up one electron and becomes a neutral chlorine atom, which is oxidation. Two such atoms then pair off into a molecule of chlorine gas.
\[\text{At the anode: } \mathrm{Cl^{-}} - e^{-} \rightarrow \mathrm{Cl} \quad \text{(oxidation)}\] \[\mathrm{Cl} + \mathrm{Cl} \rightarrow \mathrm{Cl_2}\]
Adding the two halves together gives the overall change.
\[2\,\mathrm{NaCl} \rightarrow 2\,\mathrm{Na} + \mathrm{Cl_2}\]
Reduction at the cathode, oxidation at the anode, and the same number of electrons at each — that pattern fits every cell in this chapter.
6.6 The Electrochemical Series
Section 6.5 said that cations are discharged at the cathode and anions at the anode. It did not say what happens when several different cations are present at once and only one of them can be discharged. To answer that, the ions have to be put in order.
Metal atoms lose electrons from the outermost shell and become positively charged metal ions. The number lost is the valency of the metal.
\[\mathrm{K} \rightarrow \mathrm{K^{+}} + e^{-}\] \[\mathrm{Ca} \rightarrow \mathrm{Ca^{2+}} + 2e^{-}\] \[\mathrm{Na} \rightarrow \mathrm{Na^{+}} + e^{-}\] \[\mathrm{Mg} \rightarrow \mathrm{Mg^{2+}} + 2e^{-}\] \[\mathrm{Al} \rightarrow \mathrm{Al^{3+}} + 3e^{-}\]
Every metal does this, but not with equal willingness. Some part with their valence electrons at the least provocation; others hold on to them tightly. Arrange the metals in decreasing order of that willingness and you have the electrochemical series, which decides almost everything in the rest of this chapter.
6.6.1 Electrochemical Series Of Metals
Sodium is converted to \(\mathrm{Na^{+}}\) far more readily than copper is converted to \(\mathrm{Cu^{2+}}\). Zinc becomes \(\mathrm{Zn^{2+}}\) more readily than silver becomes \(\mathrm{Ag^{+}}\). Comparisons of this kind, made for every metal, produce the order in Table 6.7.
Table 6.7 — Electrochemical Series Of Metals
| Metal | Cation formed |
|---|---|
| K | \(\mathrm{K^{+}}\) |
| Ca | \(\mathrm{Ca^{2+}}\) |
| Na | \(\mathrm{Na^{+}}\) |
| Mg | \(\mathrm{Mg^{2+}}\) |
| Al | \(\mathrm{Al^{3+}}\) |
| Zn | \(\mathrm{Zn^{2+}}\) |
| Fe | \(\mathrm{Fe^{3+}}\) |
| Ni | \(\mathrm{Ni^{2+}}\) |
| Sn | \(\mathrm{Sn^{4+}}\) |
| Pb | \(\mathrm{Pb^{2+}}\) |
| [H] | \(\mathrm{H^{+}}\) |
| Cu | \(\mathrm{Cu^{2+}}\) |
| Hg | \(\mathrm{Hg^{2+}}\) |
| Ag | \(\mathrm{Ag^{+}}\) |
| Au | \(\mathrm{Au^{3+}}\) |
| Pt | \(\mathrm{Pt^{4+}}\) |
Read it downwards and the ease of losing electrons falls away; read it upwards and the ease of gaining them back does the same. The cations at the foot of the table are therefore the ones discharged most readily at a cathode, as Fig. 6.6-1 sets out.
Diagram
Fig. 6.6-1 — The Two Series And Their Directions Two tall narrow vertical bars side by side. Left bar headed ‘Metals’, listing from top to bottom: K, Ca, Na, Mg, Al, Zn, Fe, Ni, Sn, Pb, [H], Cu, Hg, Ag, Au, Pt; an arrow beside it pointing downwards labelled ‘increasing ease of discharge at the cathode’ and a second arrow pointing upwards labelled ‘increasing ease of losing electrons to form ions’. Right bar headed ‘Anions’, listing from top to bottom: SO₄²⁻, NO₃⁻, Cl⁻, Br⁻, I⁻, OH⁻; an arrow beside it pointing downwards labelled ‘increasing ease of discharge at the anode’.
What The Series Tells You
- A metal higher in the series displaces one that is lower from a solution of its salt. The higher metal gives up its electrons more readily, and the lower metal’s ions take them.
- Metals below hydrogen cannot displace hydrogen from acids, while those above it can. Hydrogen’s place in the table is a dividing line.
- Elements lower in the series are discharged more easily at the cathode during electrolysis, because their cations gain electrons easily.
- The metal atoms that form their ions most easily accept the electrons back with the greatest difficulty. A metal cannot be both eager to lose electrons and eager to take them back.
- Hydrogen shows some properties of a metal — it forms a positive ion, for one — and it has therefore been placed among the metals in the series.
6.6.2 Electrochemical Series Of Anions
Anions are negatively charged, so on electrolysis they move to the anode, where they lose their extra electrons and are discharged. They do not all lose them with the same readiness, so they too can be arranged in order, this time of increasing ease of losing electrons.
Table 6.8 — Electrochemical Series Of Anions
| Anion | Position |
|---|---|
| \(\mathrm{SO_4^{2-}}\) | highest |
| \(\mathrm{NO_3^{-}}\) | |
| \(\mathrm{Cl^{-}}\) | |
| \(\mathrm{Br^{-}}\) | |
| \(\mathrm{I^{-}}\) | |
| \(\mathrm{OH^{-}}\) | lowest |
Two rules follow from it, and they are the mirror image of the rules for metals.
- The lower the position of an anion in the series, the more easily it is discharged at the anode.
- The higher the position of an anion in the series, the more difficult it is for that anion to lose electrons and be oxidised.
The hydroxide ion sits at the foot of the list, and the sulphate ion at the head. That single fact will explain, in Section 6.8, why oxygen and not sulphur dioxide comes off the anode when copper sulphate solution is electrolysed with inert electrodes.
6.7 Preferential Or Selective Discharge Of Ions At Electrodes
Any aqueous electrolyte contains at least four kinds of ion, not two. The compound supplies its own cation and anion, and the water supplies hydrogen ions and hydroxide ions of its own. So two positive ions arrive at the cathode and two negative ions arrive at the anode, and they cannot all be discharged. One of each pair wins, and the choosing is called selective discharge.
Which ion wins is decided by three factors:
- the relative position of the ions in the electrochemical series;
- the relative concentration of the ions;
- the nature of the electrodes.
6.7.1 Relative Position Of The Ions In The Electrochemical Series
If all other factors are held constant, the ion that stands lower in the electrochemical series is discharged at the relevant electrode in preference to those above it.
Copper sulphate solution shows this at both electrodes at once. Begin by writing down what is in the liquid.
\[\mathrm{CuSO_4} \rightarrow \mathrm{Cu^{2+}} + \mathrm{SO_4^{2-}}\] \[\mathrm{H_2O} \rightleftharpoons \mathrm{H^{+}} + \mathrm{OH^{-}}\]
Cations present: \(\mathrm{Cu^{2+}}\) and \(\mathrm{H^{+}}\). Anions present: \(\mathrm{SO_4^{2-}}\) and \(\mathrm{OH^{-}}\).
At the cathode, both cations arrive, but copper lies below hydrogen in Table 6.7, so it is the copper ions that take electrons.
\[\mathrm{Cu^{2+}} + 2e^{-} \rightarrow \mathrm{Cu}\]
At the anode, both anions arrive, and the hydroxide ion lies below the sulphate ion in Table 6.8, so it is the hydroxide ions that give up electrons. They do so in three stages.
\[\mathrm{OH^{-}} - e^{-} \rightarrow \mathrm{OH}\] \[\mathrm{OH} + \mathrm{OH} \rightarrow \mathrm{H_2O} + \mathrm{O}\] \[\mathrm{O} + \mathrm{O} \rightarrow \mathrm{O_2}\]
The sulphate ion is never discharged in this cell. It arrives at the anode, waits, and stays as it is. The full account of this electrolysis, with its apparatus and observations, comes in Section 6.8.3.
6.7.2 Concentration Of Ions In The Electrolyte
Position is not the whole story. If one ion is present in far greater numbers than another, sheer weight of numbers can override the order of the series — an ion may then be discharged in preference to ions that lie lower than it.
Concentrated sodium chloride solution is the standard example. Again, list the ions first.
\[\mathrm{NaCl} \rightarrow \mathrm{Na^{+}} + \mathrm{Cl^{-}}\] \[\mathrm{H_2O} \rightleftharpoons \mathrm{H^{+}} + \mathrm{OH^{-}}\]
At the cathode, sodium ions and hydrogen ions both arrive, and sodium ions greatly outnumber hydrogen ions. Even so, it is the hydrogen ions that are discharged, because sodium stands far above hydrogen in the series and its ions take electrons back only with great difficulty.
\[\mathrm{H^{+}} + e^{-} \rightarrow \mathrm{H}\] \[\mathrm{H} + \mathrm{H} \rightarrow \mathrm{H_2}\]
Overall, hydrogen gas bubbles off at the cathode.
\[2\,\mathrm{H^{+}} + 2e^{-} \rightarrow \mathrm{H_2}\]
At the anode the result goes the other way. Chloride ions and hydroxide ions both arrive, and here the difference in position is small while the difference in numbers is enormous — the solution is concentrated in chloride and only faintly supplied with hydroxide. So chloride is discharged, even though hydroxide occupies the lower position.
\[\mathrm{Cl^{-}} - e^{-} \rightarrow \mathrm{Cl}\] \[\mathrm{Cl} + \mathrm{Cl} \rightarrow \mathrm{Cl_2}\]
The two electrodes in this one cell show the two factors pulling against each other, and which factor wins depends on how far apart the competing ions stand in the series.
6.7.3 Nature Of The Electrodes Used
Electrodes are of two kinds, and the choice changes what is produced.
Inert electrodes, such as graphite and platinum, take no part in the electrolytic reaction. They carry current in and out and nothing more. With inert electrodes the outcome depends only on the two factors already discussed — position in the series and concentration of ions.
Active electrodes, such as copper, nickel and silver, do take part in the reaction, and the product formed is then different.
Copper sulphate solution makes the contrast plainly. Use an active copper anode and no gas appears at all: the anode itself dissolves, sending copper ions into the solution, while copper is deposited at the cathode. Replace it with an inert platinum anode and there is nothing for the anode to contribute, so hydroxide ions are discharged instead and oxygen is given off. Both cells are set out in full in Sections 6.8.3 and 6.8.4.
Table 6.9 brings the three examples together.
Compare the last row with Section 6.7.2 and the difference is instructive. Dilute sodium chloride gives oxygen at the anode, because hydroxide’s lower position decides the matter; concentrated sodium chloride gives chlorine, because there concentration takes over.
6.8 Examples Of Electrolysis
Four cells are now worked through in full. Each one is set out in the same order — the electrolyte, the electrodes, the ions present, the reaction at each electrode, and what you would actually see. Read them with Tables 6.7 and 6.8 open beside you, because every decision in them is made by position, concentration or the nature of the electrode.
6.8.1 Electrolysis Of Molten Lead Bromide
This is the simplest cell of the four, because there is no water in it and therefore no competition at either electrode.
Electrolyte. Molten lead bromide, \(\mathrm{PbBr_2}\).
Temperature. The solid must be melted and kept molten, so the crucible is held above 380 °C by continuous heating on a burner.
Electrolytic cell. A crucible made of silica.
Electrodes. Two graphite plates, which are inert, joined by copper wires to the two terminals of a battery.
Ions present. \(\mathrm{Pb^{2+}}\) and \(\mathrm{Br^{-}}\) — lead ions and bromide ions, and nothing else.
The apparatus is shown in Fig. 6.8-1.
Diagram
Fig. 6.8-1 — Electrolysis Of Molten Lead Bromide A silica crucible standing on a pipe-clay triangle on a tripod, with a Bunsen burner beneath it; the crucible holds molten lead bromide with Pb²⁺ and Br⁻ marked in the melt. Two graphite plates hang vertically into the melt from a horizontal support above, each connected by a copper wire to a battery drawn at the top with its terminals marked + and −. Label graphite cathode, graphite anode, silica crucible, molten lead bromide, burner and battery.
The Electrode Reactions
Lead ions carry two positive charges and travel to the cathode, where each takes two electrons.
\[\text{At the cathode: } \mathrm{Pb^{2+}} + 2e^{-} \rightarrow \mathrm{Pb}\]
Bromide ions travel to the anode, where each gives up its one spare electron; the neutral atoms then pair off.
\[\text{At the anode: } \mathrm{Br^{-}} - e^{-} \rightarrow \mathrm{Br}\] \[\mathrm{Br} + \mathrm{Br} \rightarrow \mathrm{Br_2}\]
The overall reaction is the decomposition of the compound into its two elements.
\[\mathrm{PbBr_2}\text{(l)} \rightarrow \mathrm{Pb}\text{(s)} + \mathrm{Br_2}\text{(g)}\]
What You Would See
- Dark reddish brown fumes of bromine come off at the anode.
- Silvery grey lead collects at the cathode.
Bromine vapour is corrosive and poisonous, so this electrolysis is done in a fume cupboard, and the hot crucible is moved only with tongs.
Two details of the apparatus are chosen for reasons worth knowing. Silica is used for the crucible because it is non-reactive, withstands a high temperature and is itself almost a non-conductor of electricity; it is heated slowly from the outside, since a sudden flame would crack it. Graphite is preferred to platinum for the anode because graphite is unaffected by the reactive bromine vapour, which would attack many other materials.
Notice also what the cell demonstrates about the solid. Solid lead bromide will not conduct at all, because its ions are held together by electrostatic attraction and are not free; the ions become free only in the fused state. That is why the crucible has to be kept hot throughout.
6.8.2 Electrolysis Of Acidified Water Using Platinum Electrodes
Electrolyte. Acidified water — water with a little dilute sulphuric acid added.
Electrodes. Platinum foils, which are inert, connected by copper wires to the two terminals of a battery.
Dilute sulphuric acid is chosen rather than dilute nitric or hydrochloric acid because it is non-volatile, while those two are volatile acids and would escape from the cell.
Why The Acid Is Needed At All
Water in the pure state consists almost entirely of molecules. It is a polar covalent compound, and it forms ions only to a very small extent — far too few to carry a useful current, which is why distilled water appeared on the list of non-electrolytes in Section 6.2.2. Adding traces of dilute sulphuric acid speeds up that ionisation without being used up itself, and for this reason the electrolysis of acidified water is regarded as an example of catalysis.
\[\mathrm{H_2O} \rightleftharpoons \mathrm{H^{+}} + \mathrm{OH^{-}}\] \[\mathrm{H_2SO_4} \rightarrow 2\,\mathrm{H^{+}} + \mathrm{SO_4^{2-}}\]
Ions present in the solution. \(\mathrm{H^{+}}\), \(\mathrm{OH^{-}}\), \(\mathrm{SO_4^{2-}}\).
Particles present in the solution. \(\mathrm{H^{+}}\), \(\mathrm{OH^{-}}\), \(\mathrm{SO_4^{2-}}\) and \(\mathrm{H_2O}\) molecules.
Diagram
Fig. 6.8-2 — Electrolysis Of Acidified Water Using Platinum Electrodes Hofmann-type voltameter: a central reservoir tube with a funnel-shaped open top, joined at the bottom to two vertical graduated side tubes closed at the top by taps. Each side tube has a platinum foil electrode sealed into its base, with wires running to a battery. Show the left tube about twice as full of collected gas as the right. Label acidified water (water + dilute H₂SO₄), platinum electrodes, cathode, anode, hydrogen (2 volumes) and oxygen (1 volume).
Reaction At The Cathode
Hydrogen ions are the only positively charged ions in the cell, so they alone migrate to the cathode. The cathode is a reservoir of electrons; each hydrogen ion takes one and becomes a neutral hydrogen atom, and the atoms pair off into molecules that come away as hydrogen gas.
\[\mathrm{H^{+}} + e^{-} \rightarrow \mathrm{H}\] \[\mathrm{H} + \mathrm{H} \rightarrow \mathrm{H_2}\]
Reaction At The Anode
Both \(\mathrm{SO_4^{2-}}\) and \(\mathrm{OH^{-}}\) migrate to the anode. Hydroxide lies lower in the electrochemical series of anions, so it is discharged preferentially.
\[\mathrm{OH^{-}} \rightarrow \mathrm{OH} + e^{-}\] \[\mathrm{OH} + \mathrm{OH} \rightarrow \mathrm{H_2O} + \mathrm{O}\] \[\mathrm{O} + \mathrm{O} \rightarrow \mathrm{O_2}\]
So the neutral hydroxide groups recombine into water, releasing oxygen, which is given off at the anode.
Removing hydroxide ions in this way disturbs the ionic equilibrium of the water, and more water ionises to restore it.
\[\mathrm{H_2O} \rightarrow \mathrm{H^{+}} + \mathrm{OH^{-}}\]
The extra hydrogen ions produced, together with the sulphate ions already gathered there, raise the concentration of sulphuric acid near the anode, while the steady discharge of hydrogen ions lowers it near the cathode. There is thus a decrease in acidity at the cathode and an increase at the anode. Taken over the cell as a whole, however, the total acidity is unchanged — no sulphuric acid is used up or made, and the overall reaction is simply the decomposition of water.
Why The Volumes Are In The Ratio 2:1
The ratio is not an accident of the apparatus. It falls out of the electron count at the two electrodes.
At the anode, four hydroxide ions must be discharged to build one molecule of oxygen, and four electrons are released in doing so.
\[4\,\mathrm{OH^{-}} \rightarrow 4\,\mathrm{OH} + 4e^{-}\] \[2\,\mathrm{OH} + 2\,\mathrm{OH} \rightarrow 2\,\mathrm{H_2O} + \mathrm{O_2}\]
Electrons cannot pile up anywhere in the circuit, so those same four electrons must be taken up at the cathode — and four electrons discharge four hydrogen ions, making two molecules of hydrogen.
\[4\,\mathrm{H^{+}} + 4e^{-} \rightarrow 2\,\mathrm{H_2}\]
Two molecules of hydrogen for every one of oxygen. Molecules can then be read as volumes, because of Avogadro’s law.
Hence the electrolysis of water gives two volumes of hydrogen for one volume of oxygen, which is exactly what the two collecting tubes in Fig. 6.8-2 show.
\[2\,\mathrm{H_2O} + \text{electrical energy} \rightarrow 2\,\mathrm{H_2} + \mathrm{O_2}\]
6.8.3 Electrolysis Of Copper Sulphate Solution Using Platinum Electrodes
Electrolyte. A saturated solution of copper(II) sulphate, prepared in distilled water with a small amount of concentrated sulphuric acid to increase the conductivity of the electrolyte.
Electrodes. Cathode of platinum, anode of platinum — both inert.
Dissociation.
\[\mathrm{CuSO_4} \rightleftharpoons \mathrm{Cu^{2+}} + \mathrm{SO_4^{2-}}\] \[\mathrm{H_2O} \rightleftharpoons \mathrm{H^{+}} + \mathrm{OH^{-}}\]
Particles present. \(\mathrm{Cu^{2+}}\), \(\mathrm{H^{+}}\), \(\mathrm{SO_4^{2-}}\), \(\mathrm{OH^{-}}\) and \(\mathrm{H_2O}\) molecules.
Ions present. \(\mathrm{Cu^{2+}}\), \(\mathrm{H^{+}}\), \(\mathrm{SO_4^{2-}}\) and \(\mathrm{OH^{-}}\).
Diagram
Fig. 6.8-3 — Electrolysis Of Copper Sulphate Solution With Platinum Electrodes Rectangular glass cell containing blue copper sulphate solution, with two flat platinum plates hanging into it and joined by wires to a battery and a rheostat drawn above. Mark the plate on the left as the platinum anode with a + beside its terminal and show small bubbles rising from it; mark the plate on the right as the platinum or copper cathode with a − beside its terminal and show a brown deposit forming on it. Label copper sulphate solution (saturated and acidified), platinum anode, platinum or copper cathode, battery.
Reaction At The Cathode
Both \(\mathrm{Cu^{2+}}\) and \(\mathrm{H^{+}}\) migrate to the cathode. Copper lies below hydrogen in the electrochemical series, so copper ions are discharged in preference to hydrogen ions and become neutral copper atoms.
\[\mathrm{Cu^{2+}} + 2e^{-} \rightarrow \mathrm{Cu}\]
These atoms deposit themselves on the cathode as a layer of pink or reddish brown copper.
Reaction At The Anode
Both \(\mathrm{SO_4^{2-}}\) and \(\mathrm{OH^{-}}\) migrate to the anode, and the hydroxide ions, standing lower in the series, are discharged to give neutral hydroxide groups. These reunite to form water and oxygen.
\[4\,\mathrm{OH^{-}} - 4e^{-} \rightarrow 4\,\mathrm{OH}\] \[2\,\mathrm{OH} + 2\,\mathrm{OH} \rightarrow 2\,\mathrm{H_2O} + \mathrm{O_2}\]
Product at the anode. Oxygen.
Product at the cathode. Pink or reddish brown copper.
What Happens To The Blue Colour
The blue of the solution is the colour of the copper ions in it. Every ion discharged at the cathode is one ion fewer in the solution, and nothing is putting any back — the platinum anode contributes nothing. So the blue fades steadily, and once the copper ions are exhausted the solution is colourless.
At that point the cell does not stop. With no copper ions left, the hydrogen ions are next in line, and what is being electrolysed from then on is the water: hydrogen is liberated at the cathode and oxygen at the anode.
Two general rules can be read off this cell, and they hold for every cell in the chapter: the oxidation of non-metal ions always takes place at the anode, and the reduction of metal or hydrogen ions always takes place at the cathode.
6.8.4 Electrolysis Of Copper Sulphate Solution Using Copper Electrodes
Change only the electrodes, and the same electrolyte behaves quite differently.
Electrolyte. Aqueous copper sulphate, acidified with traces of dilute sulphuric acid to enhance the conductivity of the electrolyte.
Electrodes. Cathode of copper, anode of copper — both active.
Dissociation.
\[\mathrm{CuSO_4} \rightleftharpoons \mathrm{Cu^{2+}} + \mathrm{SO_4^{2-}}\] \[\mathrm{H_2O} \rightleftharpoons \mathrm{H^{+}} + \mathrm{OH^{-}}\]
Ions present. \(\mathrm{Cu^{2+}}\), \(\mathrm{H^{+}}\), \(\mathrm{SO_4^{2-}}\) and \(\mathrm{OH^{-}}\) — copper ions, hydrogen ions, sulphate ions and hydroxyl ions.
Diagram
Fig. 6.8-4 — Electrolysis Of Aqueous Copper Sulphate With Copper Electrodes Rectangular glass cell of blue copper sulphate solution with two copper plates hanging into it, wired to a battery above with terminals marked. Draw the left-hand plate visibly thinner and eaten away at the edges and label it copper anode; draw the right-hand plate thicker with a fresh deposit and label it copper cathode. Label aqueous copper sulphate and battery, and add small arrows in the liquid showing Cu²⁺ moving from anode to cathode.
Reaction At The Anode
\(\mathrm{SO_4^{2-}}\) and \(\mathrm{OH^{-}}\) ions both migrate to the anode, but neither of them is discharged. The reason is the nature of the anode: it is an active electrode, and copper loses electrons more easily than either of those ions does. So the anode itself is oxidised, passing into solution as copper ions.
\[\mathrm{Cu} - 2e^{-} \rightarrow \mathrm{Cu^{2+}}\]
Reaction At The Cathode
Both \(\mathrm{Cu^{2+}}\) and \(\mathrm{H^{+}}\) ions migrate to the cathode, and copper, being lower in the series, is discharged as neutral copper atoms.
\[\mathrm{Cu^{2+}} + 2e^{-} \rightarrow \mathrm{Cu}\]
Product at the anode. No gas at all. The anode keeps dissolving as copper ions are formed and pass into the solution.
Product at the cathode. Pink or reddish brown copper, deposited on the plate.
What You Would See
The blue colour of the solution stays exactly as it was. For every copper ion discharged at the cathode as a neutral atom, one copper ion is released into the solution by the anode, so the number of copper ions never changes and the colour never fades. Over time the cathode grows thicker and the anode grows thinner as it dissolves away.
The hydrogen ions, sulphate ions and hydroxide ions in this cell arrive at the electrodes, take no part in the reaction whatever, and leave as they came. Ions of this kind are called spectator ions.
6.9 Applications Of Electrolysis
Everything in this chapter so far has been done in a beaker. The same reactions, scaled up, do three jobs of real commercial importance:
- electroplating one metal with another;
- electrolytic refining of metals, that is purifying them;
- extracting metals from their compounds, which is called electrometallurgy.
All three depend on a single fact met in Section 6.8.4 — that an active anode dissolves while the cathode gains.
6.9.1 Electroplating
Why An Article Is Plated
- For decoration. Brass objects are frequently plated with silver to give them the shine and beauty of a silver article at a fraction of its cost.
- To protect from rusting and corrosion. Iron tools and instruments are often electroplated with nickel, chromium or zinc, so that air and moisture never reach the iron beneath.
Conditions For Electroplating
Five conditions have to be met, and each of them follows from something already established.
Table 6.10 — Conditions For Electroplating And The Reason For Each
| Condition | Reason |
|---|---|
| The article to be electroplated is always placed at the cathode. | During the electrolytic reaction the metal is always deposited at the cathode, by gain of electrons. |
| The metal to be plated on the article is always made the anode, and has to be replaced periodically. | The metal anode continuously dissolves as ions in the solution, and so is used up. |
| The electrolyte must contain ions of the metal that is to be plated on the article — that is, it must be a suitable salt of the plating metal. | The electrolyte dissociates into ions of that metal, which migrate to the cathode and are deposited there as neutral metallic atoms. |
| A low current should be used, for a longer time. | A higher current causes uneven deposition of the metal. A longer time at a low current builds a thicker, uniform deposit. |
| Direct current must be used, not alternating current. | Alternating current makes discharge and ionisation alternate at the cathode, so no effective coating is formed. |
6.9.1.1 Electroplating An Article With Silver
Take a brass spoon as the article.
Preparing The Article
The spoon must be chemically clean, because a deposit will not hold on grease or on an oxide film. It is first washed in a hot solution of caustic soda, then in dilute hydrochloric acid, then in water; it is next washed thoroughly with a detergent and several times with distilled water, and dried. FC 6.2 sets out the whole sequence.
Flowchart
FC 6.2 — Electroplating A Brass Spoon With Silver Brass spoon → washed in hot caustic soda solution → washed in dilute hydrochloric acid → washed in water → washed with detergent and several times with distilled water → dried → made the cathode, with pure silver as the anode → direct current passed through sodium argentocyanide solution → silver deposited on the spoon → spoon removed, washed repeatedly with water → dried → polished
Electrolyte. Sodium argentocyanide, \(\mathrm{Na[Ag(CN)_2]}\), or potassium argentocyanide, \(\mathrm{K[Ag(CN)_2]}\).
Cathode. The article to be electroplated, duly cleaned.
Anode. A plate of pure clean silver.
Diagram
Fig. 6.9-1 — Electroplating A Brass Spoon With Silver Rectangular tank of electrolyte with a battery and an ammeter wired in series above it. A flat block hangs into the tank on the left, connected to the positive terminal, labelled anode (block of pure silver); a spoon hangs into the tank on the right, connected to the negative terminal, labelled cathode (article to be electroplated). Label battery, ammeter, and solution of sodium argentocyanide Na[Ag(CN)₂]; show small arrows in the liquid running from the silver block towards the spoon.
How The Electrolyte Is Prepared
Sodium silver cyanide is made by adding sodium cyanide to a solution of silver nitrate until the precipitate that first forms just dissolves again.
\[\mathrm{AgNO_3} + \mathrm{NaCN} \rightarrow \mathrm{AgCN}\downarrow + \mathrm{NaNO_3}\] \[\mathrm{AgCN} + \mathrm{NaCN} \rightarrow \mathrm{Na[Ag(CN)_2]}\]
Cyanides are violently poisonous. This preparation belongs to industry, where it is done under strict control by trained operators, and it is never attempted in a school laboratory.
Dissociation
\[\mathrm{Na[Ag(CN)_2]} \rightarrow \mathrm{Na^{+}} + \mathrm{[Ag(CN)_2]^{-}}\] \[\mathrm{[Ag(CN)_2]^{-}} \rightleftharpoons \mathrm{Ag^{+}} + 2\,\mathrm{CN^{-}}\] \[\mathrm{H_2O} \rightleftharpoons \mathrm{H^{+}} + \mathrm{OH^{-}}\]
The second of these is the important one. The complex ion releases silver ions only slowly and only a few at a time, so the solution is never rich in free \(\mathrm{Ag^{+}}\).
The Electrode Reactions
\[\text{At the cathode: } \mathrm{Ag^{+}} + e^{-} \rightarrow \mathrm{Ag} \quad \text{(from the electrolyte)}\] \[\text{At the anode: } \mathrm{Ag} - e^{-} \rightarrow \mathrm{Ag^{+}} \quad \text{(goes into the electrolyte)}\]
The negative ions, \(\mathrm{CN^{-}}\) and \(\mathrm{OH^{-}}\), migrate to the anode but none of them is discharged. Instead, because the anode is active, its own silver atoms lose electrons and pass into solution as silver ions. Those ions are attracted to the brass spoon at the cathode, where they gain electrons and become atoms of silver, depositing on the spoon as a firm coating.
Three consequences follow, and they are worth stating plainly:
- the thickness of the coating depends on how long the current is passed;
- the anode dissolves as the process goes on;
- the electrolyte keeps the same concentration throughout, so if it is coloured, the intensity of the colour does not change.
The spoon is finally taken out of the bath, washed several times with water, dried and polished.
6.9.1.2 Electroplating An Article With Nickel
Take an iron article this time. The apparatus is the same in principle, as Fig. 6.9-2 shows.
Electrolyte. Aqueous nickel sulphate, \(\mathrm{NiSO_4}\), containing a small amount of dilute sulphuric acid.
Cathode. The article to be electroplated, duly cleaned.
Anode. A block of pure nickel metal.
Diagram
Fig. 6.9-2 — Electroplating An Article With Nickel Rectangular tank of electrolyte with a battery above and an ammeter in the circuit. A flat vertical plate hangs into the tank on the left, joined to the positive terminal and labelled nickel anode; an iron article hangs into the tank on the right, joined to the negative terminal and labelled article to be plated. Label battery, ammeter to check current, and NiSO₄ solution.
Dissociation.
\[\mathrm{NiSO_4} \rightleftharpoons \mathrm{Ni^{2+}} + \mathrm{SO_4^{2-}}\] \[\mathrm{H_2O} \rightleftharpoons \mathrm{H^{+}} + \mathrm{OH^{-}}\]
Reaction at the cathode. Nickel ions are discharged as neutral nickel atoms by gain of electrons, in preference to the hydrogen ions also present, because nickel lies below hydrogen in the electrochemical series.
\[\mathrm{Ni^{2+}} + 2e^{-} \rightarrow \mathrm{Ni} \quad \text{(deposited)}\]
Reaction at the anode. \(\mathrm{SO_4^{2-}}\) and \(\mathrm{OH^{-}}\) both migrate to the anode, but neither is discharged, because the anode is active. The nickel anode itself loses electrons and supplies \(\mathrm{Ni^{2+}}\) ions to the solution.
\[\mathrm{Ni} - 2e^{-} \rightarrow \mathrm{Ni^{2+}}\]
Three rules cover every plating job, whatever the metal: the article to be plated is always kept at the cathode; the electrolyte must be a solution of a salt of the metal to be plated; and the anode is made of a strip of that same metal. Table 6.11 sets the two examples side by side.
Iron rusts easily when exposed to air. If iron articles are coated with a metal such as nickel, tin or zinc, rusting can be prevented.
6.9.2 Electrolytic Refining Of Copper
The most important use of copper is as wire for carrying electricity, and for that the copper must be almost perfectly pure, because very small amounts of impurity reduce its conductivity a great deal. Refining is therefore not a finishing touch but the point of the whole operation.
Electrolyte. A solution of copper sulphate and dilute sulphuric acid.
Cathode. Thin strips of pure copper.
Anode. Slabs of impure copper.
Diagram
Fig. 6.9-3 — Electrolytic Refining Of Copper Wide rectangular tank of electrolyte with many plates hanging vertically into it, alternating: thin plates labelled strips of pure copper (cathodes) and thick slabs labelled slabs of impure copper (anodes), the thin ones all wired to the negative terminal and the thick ones to the positive terminal of a battery drawn at the right. Show a dark layer settled on the tank floor beneath the anodes labelled anode mud. Label solution of copper sulphate and sulphuric acid.
Reactions.
\[\text{At the cathode: } \mathrm{Cu^{2+}} + 2e^{-} \rightarrow \mathrm{Cu}\] \[\text{At the anode: } \mathrm{Cu} - 2e^{-} \rightarrow \mathrm{Cu^{2+}}\]
Only copper is transferred. Copper ions leave the impure anode, cross to the cathode in a steady procession, gain electrons there and deposit as pure metal. The impurities cannot follow: the anode’s copper dissolves and its impurities do not, so gradually the impure slabs are eaten away and the thin strips of pure copper grow thicker and thicker. The copper deposited on the cathode is 99.9 per cent pure. When the electrolysis is over, these strips are melted and cast into the shapes wanted.
Some impurities dissolve in the acid. Others — notably silver and gold, which are insoluble — simply fall and collect below the anode as anode mud, from which they are recovered. In many refineries the value of the silver and gold pulled out of that sludge pays for a good part of the operation.
Zinc, lead, mercury, silver and copper are all refined in this way.
6.9.3 Electrometallurgy
Electrometallurgy is the extraction of metals by electrolysis.
The metals high in the electrochemical series — potassium, calcium, sodium, magnesium and aluminium — cannot be got out of their compounds by ordinary chemical means. They have a strong affinity for oxygen, so their oxides are highly stable, and common reducing agents such as carbon monoxide and hydrogen will not reduce them. Electrolysis provides a convenient method where chemistry alone fails. This was exactly Davy’s difficulty, and exactly how he solved it.
These reactive metals are extracted from their halides — chlorides, bromides and the like — by electrolytic reduction, using inert electrodes. The cathode is usually made of iron and the anode of graphite.
Why The Compound Must Be Fused, Not Dissolved
A reactive metal cannot be extracted from an aqueous solution of its salt. The solution would contain hydrogen ions from the water alongside the metal ions, and since the metal stands far above hydrogen in the series, the hydrogen ion is discharged in preference. The product at the cathode would then be hydrogen gas and not the metal at all. Fusing the salt removes the water, and with it the competition.
Four Extractions
Potassium. Electrolyte: fused potassium bromide.
\[\mathrm{KBr} \rightleftharpoons \mathrm{K^{+}} + \mathrm{Br^{-}}\] \[\text{At the cathode: } \mathrm{K^{+}} + e^{-} \rightarrow \mathrm{K}\] \[\text{At the anode: } \mathrm{Br^{-}} - e^{-} \rightarrow \mathrm{Br}; \qquad \mathrm{Br} + \mathrm{Br} \rightarrow \mathrm{Br_2}\]
Potassium metal is obtained at the cathode.
Sodium. Electrolyte: fused sodium chloride. The reactions are those already written in Section 6.5, and sodium metal is obtained at the cathode.
Magnesium. Electrolyte: fused magnesium chloride.
\[\mathrm{MgCl_2} \rightleftharpoons \mathrm{Mg^{2+}} + 2\,\mathrm{Cl^{-}}\] \[\text{At the cathode: } \mathrm{Mg^{2+}} + 2e^{-} \rightarrow \mathrm{Mg}\] \[\text{At the anode: } \mathrm{Cl^{-}} - e^{-} \rightarrow \mathrm{Cl}; \qquad \mathrm{Cl} + \mathrm{Cl} \rightarrow \mathrm{Cl_2}\]
Magnesium metal is obtained at the cathode.
Aluminium. Aluminium is obtained by the electrolysis of its oxide, alumina, \(\mathrm{Al_2O_3}\). Electrolyte: pure alumina, with cryolite, \(\mathrm{Na_3AlF_6}\), and fluorspar, \(\mathrm{CaF_2}\).
\[\mathrm{Al_2O_3} \rightleftharpoons 2\,\mathrm{Al^{3+}} + 3\,\mathrm{O^{2-}}\] \[\text{At the cathode: } 2\,\mathrm{Al^{3+}} + 6e^{-} \rightarrow 2\,\mathrm{Al}\] \[\text{At the anode: } 3\,\mathrm{O^{2-}} - 6e^{-} \rightarrow 3[\mathrm{O}]; \qquad 2[\mathrm{O}] \rightarrow \mathrm{O_2}\]
Taken together, the overall change is
\[2\,\mathrm{Al_2O_3} \rightarrow 4\,\mathrm{Al} + 3\,\mathrm{O_2}\]
6.10 Summary
Metals and alloys conduct electricity in the solid and molten states and are not changed by it; they are conductors. Non-metals other than graphite do not conduct, and are non-conductors or insulators.
Conduction in a conductor is the flow of mobile electrons. An insulator has no mobile electrons, so no current flows through it.
An electrolyte is a compound that allows an electric current to pass in the molten state or in aqueous solution and is decomposed by it. A non-electrolyte allows no current in either state and is not decomposed.
In a metallic conductor a single stream of electrons carries the charge and the change is purely physical. In an electrolytic conductor two streams of ions carry it, and new products are formed at the electrodes.
A strong electrolyte is almost completely dissociated and conducts well; a weak electrolyte is only partly dissociated and conducts poorly. The measure of this is the degree of dissociation.
\[\alpha = \frac{\text{number of molecules dissociated}}{\text{total number of molecules}} \times 100\]
A solution of a strong electrolyte contains almost only ions; a solution of a weak electrolyte contains ions as well as molecules; a non-electrolyte contains only molecules.
Acids and alkalis produce ions on dissolving in water. Bases and salts release ions already present in them, on melting or on dissolving. Salts as a family are strong electrolytes.
Electrolysis is the process in which an electric current passed through an electrolyte in its molten or aqueous state brings about a chemical reaction that decomposes it into its constituent ions.
An electrolytic cell converts electrical energy into chemical energy and needs an outside direct-current supply. An electrochemical cell does the reverse, converting chemical energy into electrical energy.
The anode is the electrode joined to the positive terminal of the battery; the cathode is the electrode joined to the negative terminal. Electrons leave the electrolyte at the anode and enter it at the cathode.
An ion is an atom or group of atoms carrying a charge equal to its valency. Cations are positive and migrate to the cathode; anions are negative and migrate to the anode.
In the electronic concept, oxidation is the loss of electrons and reduction is the gain of electrons. The older descriptions — addition of oxygen or removal of hydrogen for oxidation, and the reverse for reduction — remain valid.
An oxidising agent accepts electrons, or supplies oxygen or an electronegative ion. A reducing agent supplies electrons, or supplies hydrogen or an electropositive ion.
Arrhenius, in 1887, proposed that an electrolyte dissolved in water breaks into free mobile ions, that these ions carry the current, and that conductivity depends on their concentration.
The modern view differs from his on one point: an electrolyte is ionic even as a solid, its ions being held immobile by strong electrostatic forces. Water does not make the ions; it frees them.
Electrolytic dissociation is the separation of a compound into ions that were already present in it, and occurs with electrovalent compounds on melting or dissolving.
Ionisation is the formation of ions from polar covalent molecules that had none, and requires water. Hydrogen chloride ionises in water to give the hydronium ion.
\[\mathrm{HCl} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^{+}} + \mathrm{Cl^{-}}\]
During electrolysis, only metals and hydrogen are liberated at the cathode, and only non-metals at the anode. The process is a redox reaction — reduction at the cathode, oxidation at the anode — and it requires direct current.
Faraday’s law: the mass of a substance liberated or deposited at an electrode is directly proportional to the quantity of electricity passed through the electrolyte. One faraday is 96 500 coulombs.
The electrochemical series arranges metals in order of the ease with which their atoms lose electrons. The lower a metal stands in the series, the more readily its cation is discharged at the cathode.
For anions, the lower the position in the series, the more easily the ion is discharged at the anode. The order runs \(\mathrm{SO_4^{2-}}\), \(\mathrm{NO_3^{-}}\), \(\mathrm{Cl^{-}}\), \(\mathrm{Br^{-}}\), \(\mathrm{I^{-}}\), \(\mathrm{OH^{-}}\).
When ions of the same charge compete at one electrode, one is discharged in preference to the others. This selective discharge depends on position in the electrochemical series, on the relative concentration of the ions, and on the nature of the electrodes.
Inert electrodes such as graphite and platinum take no part in the reaction. Active electrodes such as copper, nickel and silver do take part, and the product is then different.
Molten lead bromide, electrolysed between graphite electrodes in a silica crucible, gives silvery grey lead at the cathode and reddish brown fumes of bromine at the anode.
\[\mathrm{PbBr_2}\text{(l)} \rightarrow \mathrm{Pb}\text{(s)} + \mathrm{Br_2}\text{(g)}\]
- Acidified water electrolysed between platinum electrodes gives hydrogen at the cathode and oxygen at the anode, in the ratio of two volumes to one. Dilute sulphuric acid is used because it is non-volatile, and it acts as a catalyst for the ionisation of water.
\[2\,\mathrm{H_2O} + \text{electrical energy} \rightarrow 2\,\mathrm{H_2} + \mathrm{O_2}\]
Copper sulphate solution electrolysed between platinum electrodes deposits copper at the cathode and gives oxygen at the anode, and the blue colour fades as copper ions are used up.
The same solution electrolysed between copper electrodes deposits copper at the cathode while the anode dissolves, so the blue colour does not change. The hydrogen, sulphate and hydroxide ions take no part and are spectator ions.
In electroplating, the article is the cathode, the plating metal is the anode, and the electrolyte is a salt of the plating metal. A low direct current for a long time gives a thick, uniform, adherent coat.
A brass spoon is silver-plated in sodium argentocyanide solution, which releases silver ions slowly and so deposits them smoothly. An iron article is nickel-plated in nickel sulphate solution, chiefly to prevent rusting.
In the electrolytic refining of copper, thin strips of pure copper form the cathode and slabs of impure copper the anode, in copper sulphate and sulphuric acid. The copper deposited is 99.9 per cent pure, and silver and gold collect below the anode as anode mud.
Electrometallurgy is the extraction of metals by electrolysis. Reactive metals such as potassium, sodium, magnesium and aluminium are obtained from their fused compounds, because in aqueous solution hydrogen would be discharged in their place.
The whole of the chapter’s predictive work can be reduced to the sequence in FC 6.3.
Flowchart
FC 6.3 — Predicting The Products Of Any Electrolysis Write the dissociation of the electrolyte and of water → list every cation and every anion present → check the electrodes: if the anode is active, the anode itself dissolves and no anion is discharged → if the electrodes are inert, compare the competing ions’ positions in the electrochemical series → check whether one ion is present in overwhelming concentration, which can override position → write the cathode reaction (gain of electrons, reduction) and the anode reaction (loss of electrons, oxidation) → name the product at each electrode and state what would be seen
6.11 Key Terms
Active electrodes — Electrodes such as copper, nickel and silver, which take part in the electrolytic reaction and dissolve at the anode.
Anion — A negatively charged ion, which migrates to the anode during electrolysis.
Anode — The electrode connected to the positive terminal of the battery; electrons leave the electrolyte here and oxidation takes place.
Anode mud — The insoluble residue, containing silver and gold, that collects below the anode during the electrolytic refining of copper.
Cathode — The electrode connected to the negative terminal of the battery; electrons enter the electrolyte here and reduction takes place.
Cation — A positively charged ion, which migrates to the cathode during electrolysis.
Conductors — Substances such as metals and alloys that allow an electric current to pass through them without undergoing any chemical change.
Degree of dissociation — The extent to which an electrolyte breaks up into ions, expressed as a percentage of the total number of molecules.
Electrochemical cell — A device that converts chemical energy into electrical energy, such as a simple voltaic cell.
Electrochemical series — The arrangement of metals in order of the ease with which their atoms lose electrons to form positive ions.
Electrodes — The metal plates, wires or carbon rods dipping into the electrolyte, through which the current enters and leaves the cell.
Electrolysis — The decomposition of an electrolyte, in the molten state or in aqueous solution, by the passage of direct electric current through it.
Electrolytes — Compounds that allow an electric current to pass in the molten state or in aqueous solution and are decomposed by it.
Electrolytic cell — The vessel, with its two electrodes and electrolyte, in which electrolysis is carried out; it converts electrical energy into chemical energy.
Electrolytic dissociation — The separation of an ionic compound into ions already present in it, on melting or on dissolving in water.
Electrolytic refining — The purification of an impure metal by electrolysis, using the impure metal as the anode and the pure metal as the cathode.
Electrometallurgy — The extraction of a metal from its compound by electrolysis.
Electronegative elements — The non-metals, which form negative ions and are therefore liberated at the anode.
Electroplating — The depositing of a thin film of one metal on an article made of another, with the help of electricity.
Electropositive elements — The metals and hydrogen, which form positive ions and are therefore liberated at the cathode.
Hydronium ion — The ion \(\mathrm{H_3O^{+}}\), formed when a hydrogen ion combines with a water molecule.
Inert electrodes — Electrodes such as graphite and platinum, which take no part in the electrolytic reaction.
Ion — An atom, or group of atoms, carrying a positive or negative charge equal to its valency.
Ionisation — The formation of ions from polar covalent molecules when they are dissolved in water.
Non-conductors — Substances such as most non-metals, which do not allow an electric current to pass through them.
Non-electrolytes — Compounds that allow no current to pass, either in solution or in the molten state, and are not decomposed.
Oxidation — The loss of electrons by an atom or ion; also the addition of oxygen or the removal of hydrogen.
Oxidising agent — A substance that oxidises another, by accepting electrons or by supplying oxygen or an electronegative ion.
Reducing agent — A substance that reduces another, by supplying electrons or by supplying hydrogen or an electropositive ion.
Reduction — The gain of electrons by an atom or ion; also the removal of oxygen or the addition of hydrogen.
Selective discharge — The discharge of one ion in preference to others of the same charge competing at the same electrode.
Spectator ions — Ions that are present in the electrolyte and migrate to an electrode but take no part in the electrolytic reaction.
Strong electrolyte — An electrolyte that is almost completely dissociated into ions and is therefore a good conductor.
Weak electrolyte — An electrolyte that is only partly dissociated, so that its solution contains molecules as well as ions.
6.12 Points To Ponder
Early in this chapter water appears on the list of non-electrolytes, and later on a whole cell is built to electrolyse it. Is that a contradiction, or a question of degree? What measurement would settle it?
Molten lead bromide has to be held above 380 °C before it will conduct, while copper sulphate solution conducts at room temperature. Both cells need the same thing — ions free to move. Why should one route to that freedom cost so much more energy than the other?
In one cell the electrochemical series decides which ion is discharged, and in another concentration overrules it. If both factors can win, what would you want to know about an unfamiliar solution before predicting its products?
When copper sulphate is electrolysed between copper electrodes, the anode loses exactly as much copper as the cathode gains and the solution is unchanged at the end. Nothing new has been made. On what, then, has the electrical energy been spent?
Both tin and zinc are used to coat iron against rust. Using the positions of iron, tin and zinc in the electrochemical series, would you expect a deep scratch through a tin coating and a deep scratch through a zinc coating to have the same consequences for the iron beneath?
Davy could not obtain potassium from an aqueous solution, and today potassium is still extracted from a fused compound. In two hundred years the chemistry has not been got round. Does that suggest a limit set by the apparatus, or one set by the metal itself?
6.13 Objective Questions
Multiple Choice
Choose the correct option.
- During the electrolysis of molten lead bromide, the substance formed at the anode is [1]
- lead
- bromine
- hydrogen
- oxygen
- The aqueous solution that contains ions as well as unsplit molecules is that of [1]
- sodium chloride
- acetic acid
- cane sugar
- potassium hydroxide
- An electrode that takes no part in the electrolytic reaction is [1]
- copper
- nickel
- platinum
- silver
- When dilute potassium chloride solution is electrolysed between platinum electrodes, the ion discharged at the cathode is [1]
- \(\mathrm{K^{+}}\)
- \(\mathrm{H^{+}}\)
- \(\mathrm{Cl^{-}}\)
- \(\mathrm{OH^{-}}\)
- Cryolite and fluorspar are added to the electrolyte in the extraction of [1]
- sodium
- magnesium
- aluminium
- potassium
- The copper deposited at the cathode during the electrolytic refining of copper is about [1]
- 90 per cent pure
- 95 per cent pure
- 99.9 per cent pure
- 100 per cent pure
- In the electroplating of an article with nickel, the anode is made of [1]
- the article to be plated
- pure nickel
- graphite
- platinum
- Electrons enter the electrolyte through the [1]
- anode, which is negative
- anode, which is positive
- cathode, which is negative
- cathode, which is positive
Assertion–Reason
For each pair below, choose from these four codes. (1) Both A and R are true and R is the correct explanation of A. (2) Both A and R are true but R is not the correct explanation of A. (3) A is true but R is false. (4) A is false but R is true.
Assertion (A): Solid sodium chloride does not conduct electricity. Reason (R): Solid sodium chloride contains no ions. [1]
Assertion (A): Sodium argentocyanide is preferred to silver nitrate as the electrolyte in silver plating. Reason (R): It releases silver ions slowly, so the deposit is smooth and firm. [1]
Assertion (A): Alternating current is not used to carry out electrolysis. Reason (R): The electrodes change sign repeatedly, so no effective deposit is built up. [1]
Match The Following
- Match each cell in Column A with the correct observation in Column B. [4]
| Column A | Column B |
|---|---|
| a. Molten lead bromide, graphite electrodes | i. The blue colour of the electrolyte does not change |
| b. Acidified water, platinum electrodes | ii. Reddish brown fumes at the anode |
| c. Copper sulphate solution, copper electrodes | iii. The blue colour of the electrolyte fades |
| d. Copper sulphate solution, platinum electrodes | iv. Two volumes of gas at one electrode to one volume at the other |
Fill In The Blanks
- Complete each statement. [4]
- The electrode joined to the negative terminal of the battery is the …………… .
- Anions are discharged at the anode by the …………… of electrons.
- Ions formed from polar covalent molecules dissolved in water arise by …………… .
- Ions that migrate to an electrode but are not discharged are called …………… ions.
True Or False
- State whether each is true or false. [3]
- Every substance that conducts electricity is an electrolyte.
- A weak electrolyte has a low degree of dissociation.
- Only non-metals are liberated at the cathode during electrolysis.
Name The Following
- Name the following. [3]
- A salt that behaves as a weak electrolyte.
- The electrolyte used for plating an article with silver.
- The material used for the crucible in the electrolysis of molten lead bromide.
Give One Word Or Phrase
- Give one word or phrase for each. [3]
- The purification of an impure metal by electrolysis.
- The extraction of a metal from its compound by electrolysis.
- The discharge of one ion in preference to another of the same charge at the same electrode.
Odd One Out
- Pick the odd one out in each set and give a reason. [2]
- Kerosene, urea, dilute nitric acid, benzene.
- Copper, nickel, silver, graphite.
6.14 Check Your Concepts
Define electrolysis, and state which kind of current is used to bring it about. [2]
Give a reason: a graphite anode is preferred to a platinum one in the electrolysis of molten lead bromide. [2]
Why must the silica crucible in that experiment be heated slowly and from the outside? [2]
Copper is an excellent conductor of electricity, yet it is not an electrolyte. Explain. [3]
Differentiate between an inert electrode and an active electrode, stating clearly the basis of the difference and giving one example of each. [3]
Why is a little dilute sulphuric acid added to water before it is electrolysed, and why is dilute hydrochloric acid not used for the purpose? [3]
Differentiate between ionisation and dissociation, stating the basis of difference and giving one example of each. [4]
Magnesium is extracted from fused magnesium chloride and not from its aqueous solution. What would happen at the cathode if the aqueous solution were used, and why? [4]
Selective discharge at an electrode is governed by three factors. Name them, and explain with one example how concentration can override position in the electrochemical series. [5]
State any four conditions that must be observed while electroplating an article, giving the reason for each. [5]
6.15 Equations, Conversions & Identification
For the electrolysis of molten lead bromide, write the equation for the dissociation of the electrolyte, the reaction at the cathode, the reaction at the anode, and the overall reaction. [3]
Complete and balance each of the following, and state in each case whether it is oxidation or reduction. [3]
- \(\mathrm{Ag} - e^{-} \rightarrow\) ……………
- \(\mathrm{Ni^{2+}} + 2e^{-} \rightarrow\) ……………
- \(\mathrm{Br^{-}} - e^{-} \rightarrow\) ……………
Acidified water is electrolysed between platinum electrodes. Name the gas collected at each electrode, state the ratio of their volumes, and state the law that allows molecules to be read as volumes. [3]
How would you obtain each of the following? Give the electrolyte, the electrodes and the electrode reaction in each case. [4]
- Pure copper from impure copper.
- Copper ions in solution, starting from copper metal.
A molten compound is electrolysed between graphite electrodes. A silvery grey metal collects at the cathode and dense reddish brown fumes come off at the anode. Identify the compound and both products, and write the reaction at the anode. [3]
Look at Fig. 6.8-1. Name the two electrodes and the material of the vessel, and state the purpose of the burner. [3]
Look at Fig. 6.8-4. State what you would observe at the anode, at the cathode, and in the electrolyte, and give a reason for the observation about the electrolyte. [4]
Look at Fig. 6.9-1. Name the electrolyte and the anode material, and write the reaction taking place at each electrode. [3]
Dilute copper sulphate solution is electrolysed between platinum electrodes for a long time. State two changes you would observe as the electrolysis proceeds, and name the products formed at the two electrodes once all the copper ions have been discharged. [4]