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What is a redox reaction?
Oxidation is the addition of oxygen to a substance and
reduction is the removal of oxygen from a substance.
Which substances are oxidized and reduced in this reaction?
oxygen removed
reduction
lead
oxide
+
carbon  lead
carbon
+ monoxide
oxygen added
oxidation
Reduction and oxidation always take place together.
Why is this type of reaction called a redox reaction?
redox = reduction and oxidation
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Redox reactants – oxidized or reduced?
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Redox and electrons
Magnesium burns in oxygen
to form magnesium oxide.
It is obvious that the
magnesium has been
oxidized, but what has
happened to the oxygen?
A redox reaction can also
be explained in terms of the
gain or loss of electrons.
What happens to the atoms and electrons in this reaction?
magnesium + oxygen  magnesium oxide
2Mg(s)
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+
O2(g)

2MgO(s)
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Oxidation and electron loss
When magnesium burns in oxygen to form magnesium oxide,
what happens to magnesium and its electrons?
oxidized
(electrons lost)
Mg
+
O

Mg2+
O2-
 The magnesium has been oxidized.
 The Mg atom has lost 2 electrons to form a Mg2+ ion.
Oxidation is the loss of electrons.
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Reduction and electron gain
When magnesium burns in oxygen to form magnesium oxide,
what happens to oxygen and its electrons?
reduced
(electrons gained)
Mg
+
O

Mg2+
O2-
 The oxygen has been reduced.
 The O atom has gained 2 electrons to form a O2- ion.
Reduction is the gain of electrons.
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Redox and OILRIG
An easy way to remember what happens to the electrons
during oxidation and reduction is to think… OILRIG!
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Using OILRIG
What does OILRIG stand for in terms of redox reactions?
Oxidation
Is
Loss of electrons
Reduction
Is
Gain of electrons
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What is a half-equation?
Redox reactions involve the transfer of electrons.
Equations written to show what happens to the electrons
during oxidation and reduction are called half-equations.
What are the half-equations for the oxidation and reduction
processes in this reaction?
magnesium + oxygen  magnesium oxide
2Mg (s)
+
O2 (g)

2MgO (s)
oxidation: Mg  Mg2+ + 2ereduction: O2 + 4e-  2O210 of 50
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What does each half-equation show?
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Redox reactions – summary
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What are ionic compounds?
Ionic compounds are made up of positive metal ions and
negative non-metal ions. What ions are in sodium chloride?
positive
sodium ions
negative
chloride ions
The positive and negative ions in an ionic compound attract
each other strongly. It takes a lot of energy to separate them.
How does structure affect the properties of ionic compounds?
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Properties of ionic compounds
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Do molten ionic compounds conduct?
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Do dissolved ionic compounds conduct?
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Does it conduct electricity?
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Electrolysis animation
In the solid state the ions are closely packed and immobile.
There is therefore no possibility of any movement of charge
and so no current flows.
A
+
_
_ +_ +_ +_ +_+_+_
+_ +_ +_ +_ +_+_+
_ +_ +_ +_ +_+_+_
+_ +_ +_ +_ +_+_+
_ +_ +_ +_ +_+_+_
+_ +_ +_ +_ +_+_+
_ +_ +_ +_ +_+_+_
+_ +_ +_ +_ +_+_+
0
Electrolysis of sodium chloride
Task - Describe the movement of the particles if
solid sodium chloride is heated to a temperature
above its melting point.
The ions become free to move
around – just like the particles
in any liquid.
With the power supply
switched off the ions move
randomly around the whole
volume of the molten sodium
chloride.
Only a few ions are shown here for clarity – there would be
much less space between the ions in a real liquid
A
+
_
_ + _ + _ + _ + _ + _ _+ _
_+
+ _ ++ _ + _ + _ + _ + +
_ + __ + _ +_ _ + _ + _ _+ _
_ ++
+__ ++ _ + _ + _ + _ + +
_ + __ + _ + _ + +_ + _ _
_
+
_ + __ + _ + _ +
+ _ ++ _ + +
_+_+
_ _ + _ + _ + +_ +
+_
_ + __ +__ +
+ __+ _ ++_ + +
HEAT
0
Electrolysis of sodium chloride
Task – think, pair, share
1. What do you think will happen to the movement of the ions
when the power supply is switched on?
2. Will the positive ions behave differently to the negative
ions?
when the power supply is switched on the ions start to move
in an organised manner;
• the positive ions move towards the negative electrode (cathode)
• the negative ions move towards the positive electrode (anode).
Why do you think +ve ions are called CATIONS and –ve
ions are called ANIONS?
Current flows as ions start to move to electrodes.
A
+
+
_
+
_
_
0
Electrolysis of sodium chloride
1. When the negative chloride ions arrive at the anode they
each lose an electron to form a chlorine atom.
2. The chlorine atoms combine in pairs as chlorine molecules.
2Cl-  Cl2 + 2e-
Electrolysis of sodium chloride
Task – think, pair, share
1. What do you think might happen to the sodium ions when
they reach the cathode?
The sodium ions gain electrons to form sodium metal.
Na+ + e-  Na
Task – think, pair, share
1. Predict what you might observe at each electrode.
Chlorine atoms bond in pairs as chlorine molecules and
chlorine gas escapes from the cell.
A
+
Molten sodium
forms and
collects in the
bottom of the
cell.
_
0
Anode Reaction
2Cl-  Cl2 + 2e-
+
Each anion gives up an electron
which enters the external circuit
and leaves a neutral chlorine
atom.
Negatively charged chloride ions
move through the molten
electrolyte to the anode.
Cl
Cl
-
Anode Reaction
2Cl-  Cl2 + 2eThe chlorine atoms combine in
pairs to form molecules of
chlorine gas that diffuse up and
away from the electrode.
+
Cathode Reaction
Na+ + e-  Na
_
Each cation gains one
electron from the
external circuit to form a
sodium atom.
The sodium ions move to
the cathode.
Na+
At the high temperature needed to keep the electrolyte molten
the sodium will be in the liquid state and so collects at the
bottom of the cell.
Electrolysis of sodium chloride
Overall, the net result of the processes taking place in the cell
is that the compound sodium chloride has been split (lysis)
into its component elements by the electric current (electro).
2Cl-  Cl2 + 2e2Na+ + 2e-  2Na
What is electrolysis?
An ionic compound conducts electricity
when it is molten or in solution. The
current causes the ionic compound
to split up and form new substances.
This process is called electrolysis,
a word which comes from Greek
and means “splitting by electricity”.
Electrolysis has many uses, including:
 purifying copper
 plating metals with silver and gold
 extracting reactive metals, such as aluminium
 making chlorine, hydrogen and sodium hydroxide.
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What happens during electrolysis?
In electrolysis, the substance that the current passes through
and splits up is called the electrolyte.
The electrolyte contains positive and negative ions.
What happens to these ions during electrolysis?
Negative ions move
to the positive electrode
and lose electrons.
This is oxidation.
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heat
Positive ions move
to the negative electrode
and gain electrons.
This is reduction.
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Electrolysis of molten lead bromide
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Electrolysis of molten PbBr2 – redox equations
What redox processes occur at the
electrodes during the electrolysis
of molten lead bromide (PbBr2)?
At the negative electrode:
Pb2+ + 2e-  Pb (reduction)
At the positive electrode:
2Br-  Br2 + 2e- (oxidation)
What is the overall equation for the electrolysis of molten
lead bromide ?
lead bromide  lead + bromine
PbBr2 (l)
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 Pb (l) + Br2 (g)
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Electrolysis of molten PbBr2 – summary
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Why and how is aluminium extracted?
Aluminium is one of the most
useful metals in the world.
Electrolysis is used to extract
aluminium from its ore. Why
is it not possible to extract
aluminium by heating its ore
with carbon?
Aluminium ore (bauxite) has a very high
melting point (2050 °C).
For electrolysis, the ore is dissolved in
a compound called cryolite (Na3AlF6),
which lowers the melting point to 700 °C.
Why is this important economically?
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Extracting aluminium
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Extracting aluminium – redox equations
What redox processes occur at the
electrodes during the electrolysis
of aluminium oxide (Al2O3)?
At the negative electrode:
Al3+ + 3e-  Al (reduction)
At the positive electrode:
2O2-  O2 + 4e- (oxidation)
What is the overall equation for the extraction of aluminium
by electrolysis?
aluminium oxide  aluminium + oxygen
2 Al2O3 (l)
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 4 Al (l) + 3 O2 (g)
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Extracting aluminium – summary
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What are the products of electrolysis?
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Purifying copper using electrolysis
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Purifying copper – redox equations
What happens at the electrodes
during the purification of copper
by electrolysis?
At the negative electrode:
Cu2+ + 2e-  Cu (reduction)
At the positive electrode:
Cu  Cu2+ + 2e- (oxidation)
This process is carried out on a huge scale in industry and
the copper formed on the negative electrodes is 99.99% pure.
The precious metals recovered from the impurities are also
sold off and help to make this industrial process profitable.
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Purifying copper – true or false?
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Electrolysis of NaCl solution
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Products of electrolysis of NaCl solution
The electrolysis of sodium chloride solution produces
three very useful products:
 Chlorine used for killing bacteria in water,
for bleach and making plastics like PVC.
 Hydrogen used for making margarine
and fertilizers, and for rocket fuel.
 Sodium hydroxide used in many
chemical reactions, such as making soap,
neutralizing acids and making paper.
Chlorine is expected as a product of this process but
hydrogen and sodium hydroxide are surprising products.
What happens at the electrodes to form these products?
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How does the chlorine form?
In the electrolysis of NaCl solution,
the negative chloride ions
(Cl-) are attracted to the
positive electrode.
Here, the Cl- ions lose
electrons to make
chlorine atoms, which then
form chlorine molecules (Cl2).
 Are the Cl- ions oxidized or reduced? oxidized
 How many electrons are lost by each Cl- ion? one
 How many Cl- ions join to make a Cl2 molecule? two
 What is the half-equation for this redox process?
2Cl-  Cl2 + 2e- (oxidation)
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Why is sodium not formed?
In the electrolysis of sodium chloride
solution, the Na+ ions might be expected
to form sodium at the negative electrode.
Instead, hydrogen gas is produced here.
This is because sodium chloride solution
also contains H+ ions from some of the
water: H2O (l)  H+ (aq) + OH- (aq).
At the negative electrode, the H+ ions compete with the Na+
ions. The H+ ions gain electrons; the Na+ ions stay in solution.
For all ionic compounds containing a metal that is more
reactive than hydrogen, electrolysis of a solution of the
compound will produce hydrogen rather than the metal.
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How does the hydrogen form?
In the electrolysis of NaCl solution,
the positive hydrogen ions
(H+) are attracted to the
negative electrode.
Here, the H+ ions
gain electrons to make
hydrogen atoms, which then
form hydrogen molecules (H2).
 Are the H+ ions oxidized or reduced? reduced
 How many electrons are gained by each H+ ion? one
 How many H+ ions join to make a H2 molecule? two
 What is the half-equation for this redox process?
2H+ + 2e-  H2 (reduction)
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How does the sodium hydroxide form?
Sodium chloride solution has four types of ions:
 Na+ and Cl- ions from the sodium chloride
 H+ and OH- ions from the water.
The Cl- ions form chlorine at the positive
electrode and the H+ ions form hydrogen
at the negative electrode. So, what’s left?
Na+ and OH- ions are left behind and so a
solution of sodium hydroxide (NaOH) is formed.
What is the overall equation for the electrolysis of
a sodium chloride solution?
2NaCl (aq) + 2H2O (l)  H2 (g) + Cl2 (g) + 2NaOH (aq)
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Electrolysis of NaCl solution
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increasing reactivity
Hydrogen or metal?
potassium
sodium
calcium
magnesium
aluminium
(carbon)
zinc
iron
lead
(hydrogen)
copper
silver
gold
platinum
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If an ionic compound contains a metal
that is more reactive than hydrogen,
electrolysis of a solution of the compound
will produce hydrogen, not the metal.
Complete the table for these compounds.
Ionic compound
Product at the
negative electrode
potassium chloride
hydrogen
copper sulphate
copper
sodium bromide
hydrogen
silver nitrate
silver
zinc chloride
hydrogen
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Electrolysis of dilute sulfuric acid
Electrolysis can be used to split water (H2O)
into its elements, hydrogen and oxygen.
This is how hydrogen for fuel cells can be
made and how oxygen can be produced
from water on spacecraft.
Water is a covalent compound and so is
a poor conductor of electricity. However,
it does contain a few free H+ and OH- ions:
H2O (l)  H+ (aq) + OH- (aq).
The conductivity of water can be improved by adding dilute
sulfuric acid. This releases more ions so that more current
flows during electrolysis, which creates hydrogen and oxygen.
Which product will form at each electrode?
Electrolysis of dilute sulfuric acid
Electrolysis of dilute H2SO4 – redox equations
What happens at the electrodes during
the electrolysis of dilute sulfuric acid?
At the negative electrode:
2H+ + 2e-  H2 (reduction)
At the positive electrode:
4OH-  2H2O + O2 + 4e- (oxidation)
What is the overall equation for the electrolysis of dilute
sulfuric acid?
2H2O (l)  2H2 (g) + O2 (g)
Twice as much hydrogen forms as oxygen. Why is this?
In water, there are 2 hydrogen atoms for every oxygen
atom, so the ratio by volume, of H2 to O2, is 2:1.
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Oxidized or reduced?
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What happens in electrolysis?
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Electroplating
Electroplating means using electricity to coat
one metal with another, to make it look better, or
to prevent corrosion. For example, steel car
bumpers are coated with chromium. Steel cans
are coated with tin to make tins for food. And
cheap metal jewellery is often coated with silver.
The drawing on the right shows how to
electroplate a steel jug with silver.
The jug is used as the cathode. The anode is
made of silver. The electrolyte is a solution of a
soluble silver compound, such as silver nitrate.
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At the anode The silver dissolves,
forming silver ions in solution:
Ag(s)  Ag+(aq) + eAt the cathode The silver ions are
attracted to the cathode.
There they receive electrons, forming
a coat of silver on the jug:
Ag+(aq) + e-  Ag(s)
When the layer of silver is thick
enough, the jug is removed.
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Electrolytic cell and Electrochemical Cells
• An electrochemical cell is simply a device that converts chemical
energy into electrical energy when a chemical reaction is
occurring in a cell.
•
An electrolytic cell converts electrical energy into chemical
energy.
• In an electrochemical cell the reaction occurs spontaneously at
the electrodes, while an electrolytic cell reaction is not
spontaneous at the electrodes – the reaction has to be forced by
applying an external electrical current.
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The Salt Bridge
The salt bridge completes the electrical circuit and allows
ions to flow through both half-cells.
As Zn is oxidized at the anode, Zn2+ ions are formed and
enter the solution.
Cu2+ ions leave solution to be reduced at the cathode.
The salt bridge maintains electrical neutrality by allowing
excess Zn2+ ions to enter from the anode, and excess
negative ions to enter from the cathode.
A salt bridge contains nonreacting cations and anions,
often K+ and NO3-, dissolved in a gel.
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Flow of Charge in a Voltaic Cell
Electrons flow through the wire from anode to cathode.
Zn(s) → Zn2+(aq) + 2eZn2+
Cations move through the salt
bridge from the anode solution
to the cathode solution.
Cu2+(aq) + 2e- → Cu(s)
SO42Anions move through the salt
bridge from the cathode solution
to the anode solution.
By convention, a voltaic cell is shown with the anode on
the left and the cathode on the right.
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Similarities between electrolytic and electrochemical cells
•redox reaction.
•Anode: oxidation
•Cathode: reduction
•Electrons flow from anode to cathode in the external
circuit
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Differences
Electrolytic Cell
(Electrolysis)
Chemical Cell / Voltaic
Cell
Structure
With electrical supply.
No electrical supply.
Electrodes
Can be the same or
difference metal
(graphite or platinum).
Must be two different
metals.
Flows of
electrons
From anode to cathode
through external circuit.
From more
electropositive metal to
less electropositive metal
through external circuit.
Transformation
of energy
Electrical energy to
chemical energy.
Chemical energy to
electrical energy.
At positive
terminal
At anode Oxidation
occurs. Anions release
electrons at the anode.
At cathode Reduction
occurs.
Oxidizing agent gain
electrons.
At negative
terminal
At cathode Reduction
occurs.
Cations gain electrons
from the cathode.
At anode Oxidation
occurs.
Reducing agent releases
electrons
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The oxidation of Zn(s) into Zn2+ and the reduction of Cu2+
to Cu(s) occur spontaneously. In other words, the redox
reaction between Zn and Cu2+ is spontaneous.
This is caused by the difference in potential energy
between the two substances.
The difference in potential energy between the anode
and cathode dictates the direction of electrons
movement.
Electrons move from area of higher potential energy to
area of lower potential energy. In this case, the anode
has a higher potential energy so electrons move from
anode to cathode. The potential difference between the
two electrodes is measured in units of volts.
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Figure 21.6
Measuring the standard cell potential of a zinccopper cell.
The standard cell potential is designated E°cell and is
measured at a specified temperature with no current
flowing and all components in their standard states.
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Standard Electrode Potentials
The standard electrode potential (E°half-cell) is the potential
of a given half-reaction when all components are in their
standard states.
By convention, all standard electrode potentials refer to
the half-reaction written as a reduction.
The standard cell potential depends on the difference
between the abilities of the two electrodes to act as
reducing agents.
E°cell = E°cathode (reduction) - E°anode (oxidation)
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The Standard Hydrogen Electrode
Half-cell potentials are measured relative to a standard
reference half-cell.
The standard hydrogen electrode has a standard
electrode potential defined as zero (E°reference = 0.00 V).
This standard electrode consists of a Pt electrode with H2
gas at 1 atm bubbling through it. The Pt electrode is
immersed in 1 M strong acid.
2H+(aq; 1 M) + 2e-
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H2(g; 1 atm)
E°ref = 0.00V
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Figure 21.7
Determining an unknown E°half-cell with the standard
reference (hydrogen) electrode.
Oxidation half-reaction
Zn(s) → Zn2+(aq) + 2e−
Reduction half-reaction
2H3O+(aq) + 2e- → H2(g) + 2H2O(l)
Overall (cell) reaction
Zn(s) + 2H3O+(aq) → Zn2+(aq) + H2(g) + 2H2O(l)
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For standard Electrode potential,
• temperature of the reaction is assumed to be at
25oC,
• the concentration of the reactants and products
is 1M and
•
reaction occurs at 1 atm pressure.
The standard cell potential is denoted Eocell. For
Voltaic cells it will be
Eocell=Eo(cathode) - Eo(Anode)
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Example 1: Calculate Eocell for the following redox
reaction:
2Al(s) + 3Sn2+(aq) → 2Al3+(aq) + 3Sn(s)
Solution:
Oxidation: {Al(s) → Al3+(aq) +3e–} x 2
-Eo = +1.676V
Reduction: {Sn2+(aq) +2e– → Sn(s)} x 3
Eo = -0.137V
Net: 2Al(s) + 3Sn2+(aq) → 2Al3+(aq) + 3Sn(s)
Eocell = -0.137V – (-1.676V)
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Glossary
 electrode – A solid conductor of electricity, which is used
to make electrical contact with an electrolyte.
 electrolysis – The process which uses electricity to split
up compounds.
 electrolyte – A substance which conducts electricity and
can be split up by a current when molten or in solution.
 ions – Charged particles formed when atoms lose or gain
electrons.
 oxidation – A type of reaction involving the gain of oxygen
or the loss of electrons.
 redox – A type of reaction in which oxidation and reduction
take place at the same time.
 reduction – A type of reaction involving the loss of oxygen
or the gain of electrons.
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Multiple-choice quiz
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