TOPIC 1: OXIDATION AND REDUCTION

Learning outcomes
❖ Understand the processes of oxidation and reduction cathode and their importance in
the Chemical Industry.
❖ Explain redox reactions in terms of electron transfer.
❖ Understand the changes that take place during the electrolysis of some compounds.
Key words: Electrolyte, electrolytic cell, oxidation, oxidation number, oxidizing agent, redox,
reducing agent, reduction, Anode, Cathode, Electrode
Competence: You should understand oxidation and reduction in terms of gain or loss of
oxygen and in terms of electron transfer and appreciate that two processes always occur
together
Sub – topics:
❖ Oxidation in terms of oxygen, hydrogen and electron transfer
❖ Reductions in terms of oxygen, hydrogen and electron transfer
❖ Redox reactions
❖ Reducing and oxidizing agent
❖ Electrolysis and electrochemical cells
1.0 Introduction
When metal objects are left in moist air, they turn reddish-brown due to reactions with air.
These reactions are called oxidation and reduction, and they involve the transfer of oxygen,
hydrogen, or electrons. When these reactions occur together, it's called a redox reaction. In
this chapter, you'll learn about redox reactions and how oxidation and reduction always
happen together.
1.1 Oxidation in terms of oxygen, hydrogen, electron transfer
An iron nail left in moist air turns reddish-brown due to oxidation, while sliced apples exposed to air
turn brown due to oxidation as well.
We say that a substance is oxidized if it gains oxygen, or loses hydrogen, or loses electrons.

❖ Oxidation as gain of oxygen

❖ Oxidation is loss of hydrogen
❖ Oxidation is loss of electrons

1
Many laboratory experiments can be used to illustrate oxidation reactions as shown in some
examples below

Activity 1.1.1: Illustration of oxidation in terms of oxygen

In activity 1.1, we prepare a 3 cm-long magnesium ribbon by
cleaning it with sandpaper, then ignite it in a flame while holding
it with tongs. Observe the reaction, which will produce a bright
white light and a white powder/white ash substance. Collect the
resulting product on a watch glass for analysis.
Questions
(i) Explain the reaction
(ii)Write a word equation and balanced molecular equation

Activity 1.1.2: Illustration of oxidation in terms of hydrogen

For this experiment, set up the apparatus as shown in Figure 1.3,
connecting it to a source of dry ammonia gas. Heat copper (II)
oxide in a combustion tube for several minutes and pass the dry
ammonia gas over it. Observe the changes in the combustion tube
as the reaction occurs. The black solid of copper (II) oxide turned
to brown solid of copper metal, colourless liquid also was observed.
Questions
(i) Explain the reaction
(ii) Write a word equation and balanced molecular equation
(iii) Identify the reactant substance that was:
(a) oxidized (b) reduced.

Activity 1.1.3 Illustration of oxidation in terms of electron transfer

2
Sodium + chlorine sodium chloride In describing oxidation in terms of loss of
𝟐𝑵𝒂(𝒔) + 𝑪𝒍𝟐 (𝒈) 𝟐𝑵𝒂𝑪𝒍(𝒔) electrons, we use reactions not involving oxygen
Before the reaction and hydrogen for example in a reaction between
sodium metal and chlorine gas. The reaction forms
sodium chloride. In sodium chloride we have
sodium ions and chloride ions
Sodium atom has electronic configuration of 2:8:1,
whereas chlorine atom has an electronic

After the reaction configuration of 2:8:1. The two combines through

ionic bonding, where sodium has to first loose one
electron to chlorine to form sodium ion and
chlorine gains the one electron to form chloride
ion with electronic configurations of 2:8 and
2:8:8 respectively. The ions are then attracted to
form sodium chloride.
1.2 Reduction Reaction in terms of oxygen, hydrogen, electron transfer

What is reduction reaction?

A reduction reaction is a reverse process of an oxidation reaction. Reduction is said to take

place when a substance loses oxygen, gains hydrogen, or gains electrons.

Reduction reactions do not occur independently. In all oxidation reactions, there are reductions
reactions occurring simultaneously to form oxidation-reduction (RedOx) reactions. For
instance, during cellular respiration, glucose molecules are oxidized to produce energy in the form of
ATP (adenosine triphosphate), while oxygen molecules are reduced to form water. Similarly, in
photosynthesis, carbon dioxide is reduced to form glucose, while water is oxidized to release oxygen.

Many laboratory experiments can be used to illustrate oxidation reactions as shown in some
examples below

3
 Activity 1.2.1: Illustration of reduction in terms of oxygen

A mixture of zinc powder and copper (II) oxide is heated. The

two substances react according to the equation below:

In the reaction, copper (II) oxide has lost oxygen , so it has

been reduced to copper metal

Questions:

1. Suggest a possible observation

Activity 1.2.2: Illustration of reduction in terms of hydrogen

We can consider a reaction between ammonia and bromine

Ammonia gas+ bromine gas nitrogen gas+ hydrogen bromide gas

In the reaction bromine gains hydrogen to form hydrogen bromide and it is reduced.

Activity 1.2.2: Illustration of reduction in terms of electron transfer

4
 Considering the same example as in activity 1.1.3 above for a reaction between sodium metal and
chlorine gas to form sodium chloride which contains sodium ions and chloride ions

We saw that sodium atom had electronic configuration of 2:8:1, whereas chlorine atom had an
electronic configuration of 2:8:1. The two combined through ionic bonding, where sodium had to
first loose one electron to chlorine to form sodium ion and chlorine gained the one electron to form
chloride ion with electronic configurations of 2:8 and 2:8:8 respectively. The ions are then attracted
to form sodium chloride. In the reaction chlorine was reduced to form chloride ion by gaining one
electron.

Sodium + chlorine sodium chloride

𝟐𝑵𝒂(𝒔) + 𝑪𝒍𝟐 (𝒈) 𝟐𝑵𝒂𝑪𝒍(𝒔)

Reduction by gain of electrons

1.3 Importance of oxidation and reduction in everyday life

Many processes involve chemical reactions which are either oxidation or reduction reaction.

❖ Extraction of metals, like iron and zinc, involves reduction their ores.
❖ Breakdown of food by cells to produce energy is an oxidation reaction.
❖ Most of the electrochemical cells, such as car batteries, dry cells, use redox reactions
to produce electrical energy.
❖ Redox reactions are also used in electroplating which involves coating a metal with a
thin layer of another metal.

5
1.5 Redox Reactions

Oxidation and reduction always take place simultaneously. In other words, there be no
oxidation without reduction, and vice versa. We call the combined process redox reaction. A
redox reaction is a reaction in which both oxidization and reduction reactions take place at
the same time (simultaneously). This can be explained in terms of oxygen, hydrogen and
electron transfer

1. In terms of oxygen: When steam is passed over heated magnesium, magnesium oxide
and hydrogen gas are produced.

magnesium + steam magnesium oxide + hydrogen

In the reaction magnesium gained oxygen, and so has been oxidized to magnesium oxide,
while water lost oxygen and so has been has been reduced to hydrogen.

2. In terms of hydrogen: When ammonia gas is reacted with bromine gas. Ammonia
looses a hydrogen and it is oxidized while bromine gains a hydrogen and it is reduced.

3. In terms of electron transfer:

Metals lose electrons and, in the process, they are oxidized, while non-metals gain electrons
and get reduced.

When magnesium metal with electronic configuration 2:8:2 reacts with oxygen which has
electronic configuration 2:6. Magnesium loses two electrons which are transferred to oxygen
atom. Magnesium atom becomes magnesium ion while oxygen becomes the oxide ion.

6
 magnesium + oxygen magnesium oxide

There are two reaction; reduction and oxidation as below

Magnesium atoms loses electrons to form magnesium ions (oxidation reaction):

𝑀𝑔(𝑠) 𝑀𝑔2+ (𝑙) + 2𝑒

The electrons lost are gained by oxygen to form oxide ions (reduction reaction):

𝑂(𝑔) + 2𝑒 − 2𝑂2− (𝑙)

Discussion questions

1. What has been oxidized in the reaction between magnesium and oxygen?

2. What happens when copper is placed into silver nitrate solution?

Activity 1.5.1 Examples of redox reactions

You are required to work in groups.
What you need:
Internet and relevant Chemistry textbook
What to do
1. Research about examples of redox reactions and identify the oxidized or reduced species
in three (3) different reactions.
2. Record your findings in Table 1.2.
Table 1.2: Oxidized and reduced species in redox reactions
Reaction Oxidized species Reduced species

3. Write ionic equations for each of the reactions, clearly showing the electron transfers.

Discussion question

Identify the species which have been reduced in the following reaction

(a)2𝐾𝐼(𝑎𝑞) + 𝐻2 𝑂2 (𝑎𝑞) 𝐼2 (𝑠) + 2𝐾𝑂𝐻(𝑎𝑞)

7
(b)𝐹𝑒(𝑠) + 𝑆(𝑠) 𝐹𝑒𝑆(𝑠)

(d)𝐻2 𝑆(𝑔) + 𝐵𝑟2 (𝑔) 2𝐻𝐵𝑟(𝑔) + 𝑆(𝑠)

(𝑒)𝐻2 𝑆(𝑔) + 𝐶𝑙2 (𝑔) 2𝐻𝐶𝑙(𝑔) + 𝑆(𝑠)

1.6 Oxidation Number

This is the total number of electrons that an atom or ion either gains or loses in order to
form a chemical bond with another atom or ion.

Oxidation numbers can be calculated by applying the following rules:

❖ For an uncombined element in its atomic or molecular state such as, chlorine molecule
(Cl2) and zic (Zn), the oxidation number is zero.
❖ The charge on a monoatomic ion is its oxidation number, for example, the oxidation
number of the sodium ion (Na+) is +1, while that of the chloride ion (Cl-)is-1
❖ The oxygen atom in a compound or complex ion has oxidation state of -2, except in
peroxides where the oxidation state is -1.
❖ The sum of oxidation numbers of all atoms in a neutral compound is zero.
❖ The sum of oxidation numbers of all atoms in an ion or molecule made up of more
than one atom is equal to the overall charge on the ion.
❖ Hydrogen atoms have an oxidation number of +1, unless they are bonded to a metal
atom, when the oxidation number will be -1

Example 1.1

Calculate the oxidation number of nitrogen in the following substances:

a) nitrogen dioxide (NO2)

Solution:

From the chemical formula of nitrogen dioxide, NO2, which is a neutral compound, the sum
of all the oxidation numbers of the atoms it contains should be equal zero charges.

Now, let the oxidation number of nitrogen be n. Oxygen has an oxidation num of -2 and
there are two oxygen atoms in NO2.

𝑇ℎ𝑢𝑠 𝑛 + (2 × −2) = 0

8
 𝑛 + −4 = 0

n=+4.

Hence, the oxidation number of nitrogen in NO2

b) nitrate ion (NO3-)

Solution:
The last rule in Table 1.2 states that the sum of all the oxidation numbers of at in a
polyatomic ion should be equal to the overall charge on the ion.S o,let the oxidation number
of nitrogen be y.
Then𝑦 + (3𝑥 − 2) = −1
+ −6 = −1
y = +5
Thus, the oxidation number of nitrogen in the nitrate ion is +5.
Using the rules of calculating oxidation numbers you have learnt so far, you find the oxidation
numbers of different atoms in Activity 1.2(c).
Activity
Do this activity individually.
What you need:
notebook pen
What to do
1. Calculate the oxidation numbers of the following named atoms in the given compounds:
a) sulphur in SO3 d) magnesium in MgH2
b) carbon in CO32- e) nitrogen in NCl3
c) lead in PbO2 f) sodium in Na2O2

2. Present your findings to the rest of the class.

By finding the oxidation numbers of species in a reaction, you can tell what is being oxidised
or reduced. Therefore, on the basis of oxidation number, oxidation may be defined as an
increase in the oxidation number of a species in a reaction.

9
Oxidizing and Reducing Agents

You have learned that oxidation and reduction reactions take place simultaneously. In such
reactions, a substance which oxidizes another substance is called an oxidizing agent
(oxidant). Therefore, an oxidizing agent is a substance in a redox reaction that gains
electrons, or hydrogen or loses oxygen and so gets reduced in the reaction. A reducing agent
(reductant) is defined similarly. The oxidation state of an oxidizing agent reduces because
it undergoes reduction while that of the reducing agent increase because it undergoes
oxidation.

For examples:

1. In a reaction between ammonia and bromine, ammonia loses a hydrogen which is

gained by bromine

2. In a reaction between copper (II) sulpahte and zinc, the oxidation state of zinc
increases from 0 to +2 , while that of copper decreases from +2 to 0 as seen in
the ionic equation

3. In a reaction between zinc and dilute Sulphuric acid, the oxidation state of zinc
increases from 0 to +2 , while that of the hydrogens in the acid decreases from +1
to 0 as seen in the ionic equation

10
 Assignment 1.1

1. Define the term "reducing agent".

2. In each of the following reactions, identify the reducing and oxidizing agents.

(a) 𝑀𝑔(𝑠) + 𝐵𝑟2 (𝑙) 𝑀𝑔𝐵𝑟2 s)

(b) 2𝐹𝑒 3+ (𝑎𝑞) + 𝑍𝑛(𝑠) 2𝐹𝑒 2+ (𝑎𝑞) + 𝑍𝑛2+ (𝑎𝑞 2 + (𝑎𝑞)

(c)2𝐻 + (𝑎𝑞) + 𝑍𝑛(𝑠) 𝐻2 (𝑔) + 𝑍𝑛2+ (𝑎𝑞 72+ (𝑎𝑞)

(d)3𝐶𝑢𝑂(𝑠) + 2𝑁𝐻3 (𝑔) 3𝐶𝑢(𝑠) + 3𝐻2 𝑂(𝑙) + 𝑁2 (𝑔)

(e)𝐹𝑒2 𝑂3 (𝑠) + 3𝐶𝑂(𝑔) 2𝐹𝑒(𝑠) + 3𝐶𝑂2 (𝑔)

(f) 𝑀𝑛𝑂4− (𝑎𝑞) + 8𝐻 + (𝑎𝑞) + 5𝐹𝑒 2+ (𝑎𝑞) 𝑀𝑛2+ (𝑎𝑞) + 5𝐹𝑒 3+ (𝑎𝑞) + 4𝐻2

Practical examples in redox reactions

Reaction between copper and silver nitrate solution

In this experiment, a strip of copper metal is coiled and

placed into a test tube containing silver nitrate solution.
After standing for 3-4 minutes, the silver nitrate solution
becomes less clear and may turn blue due to the formation
of copper (II) nitrate, while the copper wire becomes
coated with a greyish or silver substance, indicating the
At the start of the After the reaction deposition of silver metal.

reaction Chemical Equation for the Reaction:

𝑨𝒈𝑵𝑶𝟑 (𝒂𝒒) + 𝑪𝒖(𝒔) 𝑪𝒖(𝑵𝑶𝟑 )𝟐 (𝒂𝒒) + 𝟐𝑨𝒈(𝒔)

Electron Transfer process:

Copper (Cu) loses electrons (oxidized):

𝐶𝑢(𝑠) 𝐶𝑢2+ (𝑎𝑞) + 2𝑒

11
 Silver (Ag) gains electrons (reduced):

𝐴𝑔+ (𝑎𝑞) + 𝑒 𝐴𝑔(𝑠)

Overall Ionic Equation:

𝐶𝑢(𝑠) + 𝐴𝑔+ (𝑎𝑞) 𝐶𝑢2+ (𝑎𝑞) + 2𝐴𝑔

Reason for the Reaction Being a Redox Reaction:

The reaction involves the transfer of electrons from copper

to silver ions. Copper is oxidized (loses electrons), and
silver ions are reduced (gain electrons), which defines a
redox (reduction-oxidation) process.

Effect of Sulphur dioxide gas on acidified potassium dichromate and acidified potassium permanganate

To investigate the effect of sulphur dioxide on acidified

potassium dichromate and potassium permanganate solutions,
you will need iron (II) sulphate, a retort stand, boiling tubes,
a heat source, a rubber stopper, 3cm³ of 2 M dilute sulphuric
acid, 5cm³ of potassium dichromate solution, and 5cm³ of
potassium permanganate solution. First, transfer 5cm³ of
potassium dichromate solution into a boiling tube and add 3
cm³ of dilute sulphuric acid. In another boiling tube, place 6g
of iron (II) sulphate crystals and set up the apparatus as
shown in figure, ensuring the delivery tube is dipped into the
(a) State observations made using acidified potassium dichromate solution. Heat the iron (II)
(i) Acidified potassium dichromate sulphate crystals until there is no further change. Repeat the
(ii) Acidified potassium permanganate experiment using potassium permanganate solution instead.
(b) How was Sulphur dioxide obtained during the
experiment?
(c) What was the role of Sulphur dioxide in the
experiment

Assignment: Discuss Reaction of zinc and copper (II) sulphate solution

12
1.7: Distinguishing between redox and non-redox reactions

A reaction in which neither oxidation nor reduction takes place is called a non-redox reaction.
Examples of such reactions include:

i) Solutions of barium chloride and sodium sulphate react to give insoluble barium sulphate
and the solution of sodium chloride.

BaCl2(aq)+Na2SO4(aq) BaSO4 (s)+2NaCl(aq)

ii) Sodium hydroxide solution reacts with hydrochloric acid solution to prod sodium chloride
solution and water.

𝑁𝑎𝑂𝐻(𝑎𝑞) + 𝐻𝐶𝑙(𝑎𝑞) 𝑁𝑎𝐶𝑙(𝑎𝑞) + 𝐻2 𝑂(𝑙)

Do you notice that there is neither oxidation nor reduction in the above reaction

ELECTROLYSIS
Key Terms
Anode: The electrode where oxidation occurs.
Cathode: The electrode where reduction occurs.
Conductor: A material that allows the flow of electric current.
Electrode: A conductor through which electricity enters or leaves an electrolyte.
Electrolysis: The process of using electrical energy to drive a chemical reaction. Electrolysis
can be defined the decomposition of a substance by passing an electric current through its
molten or aqueous form. This process is carried out in an electrolytic cell, which contains
an electrolyte and electrodes.
Electrolyte: A substance that conducts electricity when dissolved in water or molten.
Electrolytic Cell: A cell that uses electrical energy to drive a chemical reaction.
Oxidation: The loss of electrons.
Reduction: The gain of electrons.

Electrolysis of Some Compounds

Electrolysis can be used to decompose various compounds, typically those that are ionic
either when molten or in aqueous solution. Let's examine some examples.

13
 1.3 Electrolysis of molten compounds

Example: Electrolysis of molten Lead (II) Bromide

Materials:

Crucible, Carbon electrodes, Lead (II) bromide powder, Heat source, Connecting wires, Bulb

Procedure:

1. Place lead (II) bromide in a

crucible and heat until molten.
2. Insert carbon electrodes into
the molten lead (II) bromide.
3. Connect the electrodes to a
battery using connecting wires.
4. Observe the bulb lighting up,
indicating the conductivity of
the molten lead (II) bromide.

Observations and Analysis:

When lead(II) bromide (PbBr₂) is heated until molten, it dissociates into lead (Pb²⁺) and
bromide (Br⁻) ions. The electric current causes these ions to migrate towards the
respective electrodes.

At the Cathode (negative electrode):

Reduction occurs: Lead ions gain electrons to form lead metal.

Equation: 𝑃𝑏 2+ (𝑙) + 2𝑒 𝑃𝑏(𝑠)

Observation: Grey solid is deposited on the cathode.

At the Anode (positive electrode):

Oxidation occurs: Bromide ions lose electrons to form bromine gas.

Equation: 2𝐵𝑟 − (𝑙) 𝐵𝑟2 + 2𝑒

14
Observation: Brown gas forms at the anode, which may be observed as a reddish-brown
vapor.

Summary of Observations:

The bulb lighting up indicates that the molten lead (II) bromide conducts electricity.

Grey solid is observed at the cathode.

Reddish-brown bromine gas is released at the anode.

Assignment: Describe the electrolysis of Molten Sodium Chloride

Electrolysis of substances in aqueous form

During electrolysis in aqueous form, a number of ions are produced both from water and
the electrolyte substance for example in a solution of sodium chloride, there are sodium
ions (Na+) and chloride ions (Cl-) form sodium chloride (NaCl) and also hydrogen
ions(H+) and hydroxide ions (OH) from water. So, choice an ion for discharge at a
particular electrode depends on the following factors:

1. Position of the ion in the electrochemical series. The ion lower in

the electrochemical series is most preferred
2. Concentration of the ion in the electrolyte. Irrespective of position
of the ion in the electrochemical series, there is always a tendency
to promote to promote the discharge of the, most concentrated ion
3. Nature of the electrode

Example 1: Electrolysis of acidified water

Electrolysis of acidified water produces hydrogen gas at the cathode and oxygen gas at the
anode.

Equipment and Chemicals:

• Electrolysis apparatus
• Dilute sulphuric acid
• Power supply

15
 Procedure:

1. Fill the electrolysis apparatus with water

mixed with a small amount of sulfuric
acid or sodium sulfate to increase
conductivity.
2. Insert two electrodes into the
solution and connect them to a power
supply.
3. Turn on the power supply and
observe the formation of gas bubbles at
both electrodes.
Equations:
At the cathode: 2𝐻 + (𝑎𝑞) + 2𝑒 𝐻2 𝑂(𝑙)
At the anode: 4𝑂𝐻 − (𝑎𝑞) 𝑂2 (𝑔) + 4𝐻2 𝑂(𝑙) + 4𝑒
Observations:
• Bubbles of a colourless gas which burns with a pop sound (Hydrogen gas) forms at
the cathode.
• Bubbles of a colourless gas which relights a glowing splint (Oxygen gas) forms at
the anode
Electrolysis of Copper (II) Sulfate Solution using carbon electrodes
Electrolysis apparatus
Copper (II) sulfate solution (CuSO₄)
Copper electrodes
Power supply
Procedure:
1. Fill the electrolysis apparatus with copper
(II) sulfate solution.
2. Insert two carbon electrodes into the
solution and connect them to a power
supply.
3. Turn on the power supply and observe the
changes at both electrodes.

Equations:

16
At the cathode: 𝐶𝑢2+ (𝑎𝑞) + 2𝑒 𝐶𝑢(𝑠)

At the anode: 4𝑂𝐻 − (𝑎𝑞) 𝑂2 (𝑔) + 4𝐻2 𝑂(𝑙) + 4𝑒

Observations:

Copper metal deposits on the cathode.

Bubbles of a colourless gas which relights a glowing splint (oxygen) forms at the anode.

Note: Read about electrolysis of copper (II) sulfate (CuSO₄) solution using copper
electrodes It results in the deposition of copper on the cathode and the dissolution of
copper at the anode.
Applications of Electrolysis in the Chemical Industry
Electrolysis plays a crucial role in the chemical industry for various applications, such as:
1. Extraction of Metals: Electrolysis is used to extract metals from their ores, such as
aluminum from bauxite and sodium from sodium chloride.
2. Electroplating: Used to coat objects with a thin layer of metal to prevent corrosion,
improve appearance, or provide other benefits.
3. Production of Chemicals: Used in the production of important chemicals like chlorine,
sodium hydroxide, and hydrogen.
4. Purification of Metals: Used to purify metals like copper and zinc by removing
impurities.
5. Battery Technology: Electrolysis is involved in the operation of various batteries,
including lead-acid and lithium-ion batteries.
Electroplating
Electroplating is an application of electrolysis where a metal is deposited onto the surface
of another material.
Example: Electroplating of Iron with Copper Materials:

Iron object (cathode)

Copper electrode (anode)

Copper sulfate (CuSO₄) solution

Procedure:

17
 1. Prepare the copper sulfate solution.
2. Insert the iron object as the
cathode and the copper electrode as
the anode.
3. Connect the electrodes to a power
supply.
4. Turn on the power supply and let
the process continue until a desired
thickness of copper is plated on the
iron object.
Equations:
At the cathode: 𝐶𝑢2+ (𝑎𝑞) + 2𝑒 𝐶𝑢(𝑠)
At the anode: 𝐶𝑢(𝑠) 𝐶𝑢2+ (𝑎𝑞) + 2𝑒
Observations:
Brown solid (copper metal) deposits on the iron object (cathode).
Electrochemical cell
This a device which converts chemical energy into electrical energy. It is also called a
galvanic cell or voltaic cell.
In an electrochemical cell, the anode consisting of a more reactive metal undergoes
oxidation and is negative electrode. The cathode made of a less reactive metal undergoes
reduction and is a positive electrode. These processes result into electrical energy
Examples include car battery, simple cell, Daniel cell
Mode of operation of an electrochemical cell (Daniel cell)
The Daniel cell, named after the
British chemist John Frederic Daniel,
is an example of a galvanic (voltaic)
cell. Here's how it operates:

18
Anode Reaction:

At the anode (negative electrode), zinc (Zn) metal undergoes oxidation:

𝑍𝑛(𝑠) 𝑍𝑛2+ (𝑎𝑞) + 2𝑒

This reaction releases electrons into the external circuit.

Cathode Reaction:

At the cathode (positive electrode), copper (Cu) ions in solution gain electrons and
deposit as copper metal:

𝐶𝑢2+ (𝑎𝑞) + 2𝑒 𝐶𝑢(𝑠)

This reaction consumes electrons from the external circuit.

Overall equation (redox equation)

𝑍𝑛(𝑠) + 𝐶𝑢2+ (𝑎𝑞) 𝑍𝑛2+ (𝑎𝑞) + 𝐶𝑢(𝑠)

Electron Flow:

Electrons flow from the anode to the cathode through the external circuit, generating an
electric current that can be used to power external devices.
Salt Bridge:
A salt bridge or porous barrier separates the two half-cells, allowing ions to move between
them to maintain charge neutrality. It also completes the circuit In the Daniel cell, the
salt bridge typically contains a solution of potassium nitrate (KNO3) or other electrolytes.
Read about simple cells

19