3. Electrodes

What Is an Electrode?

An electrode is an electrical conductor that allows electrons to enter or leave an electrochemical system.

In an electrolytic cell, there are two electrodes:

Anode

and

Cathode

The electrodes are connected to an external power supply and placed in contact with an:

electrolyte.

Electrode surfaces are important because this is where:

oxidation and reduction reactions occur.

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The Two Electrodes

During electrolysis:

Anode = positive electrode

Cathode = negative electrode

The external power supply creates this difference in electrical potential.

This causes ions in the electrolyte to move toward the:

oppositely charged electrode.

Therefore:

Cations (+) → Cathode (−)

Anions (−) → Anode (+)

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The Cathode

The cathode is the electrode where:

reduction occurs.

Reduction means:

gain of electrons.

A useful memory aid is:

RedCat

REDuction at the CAThode

In an electrolytic cell, the cathode is connected to the:

negative terminal of the power supply.

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What Happens at the Cathode?

The power supply pushes electrons toward the cathode.

As a result, the cathode provides electrons to particles in the:

electrolyte.

Positive ions are attracted toward the negative cathode.

These positive ions are called:

cations.

At the electrode surface, a cation may gain one or more electrons.

For example:

Cu²⁺ + 2e⁻ → Cu

The copper ion gains two electrons and becomes:

a neutral copper atom.

This is reduction.

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The Anode

The anode is the electrode where:

oxidation occurs.

Oxidation means:

loss of electrons.

A useful memory aid is:

AnOx

ANode = OXidation

In an electrolytic cell, the anode is connected to the:

positive terminal of the power supply.

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What Happens at the Anode?

Negative ions are attracted toward the positive anode.

These negative ions are called:

anions.

At the anode, particles may lose electrons.

For example:

2Cl⁻ → Cl₂ + 2e⁻

Chloride ions lose electrons and form:

chlorine gas.

This is oxidation.

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The Essential Rules

These four rules are worth remembering:

Cathode = Reduction

Anode = Oxidation

Cations → Cathode

Anions → Anode

For an electrolytic cell:

Cathode = negative

Anode = positive

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A Useful Memory System

There are several memory aids that can help.

RedCat

Reduction occurs at the Cathode.

AnOx

Oxidation occurs at the Anode.

OIL RIG

Oxidation Is Loss of electrons.

Reduction Is Gain of electrons.

Together:

Anode → Oxidation → Loss of electrons

Cathode → Reduction → Gain of electrons

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Why Do Ions Move Toward the Electrodes?

Opposite electrical charges:

attract.

Therefore positive ions are attracted toward the:

negative cathode.

Negative ions are attracted toward the:

positive anode.

This movement of ions allows electrical charge to move through the:

electrolyte.

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Cations

A cation is a positively charged ion.

Examples include:

• Na⁺
• Cu²⁺
• Mg²⁺
• Pb²⁺
• Ag⁺

During electrolysis:

cations move toward the cathode.

At the cathode, they may:

gain electrons.

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Anions

An anion is a negatively charged ion.

Examples include:

• Cl⁻
• Br⁻
• I⁻
• OH⁻
• SO₄²⁻

During electrolysis:

anions move toward the anode.

At the anode, particles may:

lose electrons.

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Electron Movement and Ion Movement

It is important not to confuse:

electrons

with:

ions.

Electrons move through:

• wires
• electrodes
• the external circuit

Ions move through:

• molten electrolytes
• electrolyte solutions

The two forms of charge movement work together to complete the:

electrical circuit.

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Visualizing an Electrolytic Cell

The relationship between the electrodes, ion movement, electron flow, and electrode reactions is easier to understand when viewed together.

Notice the key relationships:

positive ions → cathode → reduction

negative ions → anode → oxidation

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Electrodes Are Reaction Surfaces

Electrodes do more than simply attract ions.

They provide a surface where:

electron-transfer reactions

can occur.

At the cathode, electrons are available to particles in the electrolyte.

At the anode, electrons are removed from particles.

The electrode therefore connects the:

electronic circuit

with the:

ionic system.

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Half-Equations

Electrode reactions are often represented using:

half-equations.

A half-equation shows either:

• oxidation

or:

• reduction

including the electrons involved.

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Cathode Half-Equations

A cathode half-equation shows:

gain of electrons.

For example:

Na⁺ + e⁻ → Na

The electron appears on the:

left side.

Another example:

Cu²⁺ + 2e⁻ → Cu

Copper ions gain electrons.

Therefore:

reduction has occurred.

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Anode Half-Equations

An anode half-equation shows:

loss of electrons.

For example:

2Cl⁻ → Cl₂ + 2e⁻

The electrons appear on the:

right side.

Another example:

2Br⁻ → Br₂ + 2e⁻

Bromide ions lose electrons.

Therefore:

oxidation has occurred.

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Example: Molten Sodium Chloride

Molten sodium chloride contains:

Na⁺ ions

and:

Cl⁻ ions.

Because the substance is molten, these ions are free to:

move.

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At the Cathode

Na⁺ ions move toward the:

negative cathode.

They gain electrons:

Na⁺ + e⁻ → Na

This is:

reduction.

The product is:

sodium metal.

---

At the Anode

Cl⁻ ions move toward the:

positive anode.

They lose electrons:

2Cl⁻ → Cl₂ + 2e⁻

This is:

oxidation.

The product is:

chlorine gas.

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Overall Reaction

The two half-equations can be combined.

Cathode:

2Na⁺ + 2e⁻ → 2Na

Anode:

2Cl⁻ → Cl₂ + 2e⁻

The electrons cancel.

Overall:

2NaCl → 2Na + Cl₂

Electrical energy drives this:

non-spontaneous reaction.

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Example: Molten Lead(II) Bromide

Consider molten lead(II) bromide:

PbBr₂

It contains:

Pb²⁺ ions

and:

Br⁻ ions.

The ions are mobile because the compound is:

molten.

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At the Cathode

Pb²⁺ is positive.

Therefore:

Pb²⁺ → cathode

At the cathode:

Pb²⁺ + 2e⁻ → Pb

Lead ions gain electrons.

Therefore:

reduction occurs.

Lead metal is produced.

---

At the Anode

Br⁻ is negative.

Therefore:

Br⁻ → anode

At the anode:

2Br⁻ → Br₂ + 2e⁻

Bromide ions lose electrons.

Therefore:

oxidation occurs.

Bromine is produced.

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Try the Second Electrolyte

The same electrode rules apply when the electrolyte changes.

The identities of the products change, but the basic rules remain:

Anode = oxidation

Cathode = reduction

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Inert and Reactive Electrodes

Electrodes can be classified as:

inert electrodes

or:

reactive electrodes.

This distinction is important because the electrode material can sometimes affect the products of:

electrolysis.

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Inert Electrodes

An inert electrode provides a conducting surface but does not normally participate significantly in the chemical reaction under the conditions used.

Common examples include:

• graphite
• platinum

The electrode allows electrons to enter or leave the system while remaining largely:

chemically unchanged.

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Why Use Inert Electrodes?

An inert electrode is useful when scientists want to study reactions involving the:

electrolyte itself.

The electrode provides:

• electrical contact
• a reaction surface
• electron transfer

without intentionally supplying another reactant.

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Graphite Electrodes

Graphite is commonly used as an electrode because it:

• conducts electricity
• can withstand relatively high temperatures
• is comparatively inexpensive
• is relatively unreactive in many conditions

However, graphite is not perfectly inert under:

all conditions.

It can participate in some reactions, particularly under strongly oxidizing or high-temperature conditions.

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Platinum Electrodes

Platinum is another common inert-electrode material.

It:

• conducts electricity well
• resists many chemical reactions
• provides an effective electrode surface

However, platinum is:

expensive.

Therefore graphite is often more practical in school laboratory experiments.

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Reactive Electrodes

A reactive electrode participates in the chemical reaction.

Instead of simply providing a surface, the electrode material itself may:

• lose electrons
• dissolve into the electrolyte
• gain mass
• react with products

The electrode therefore becomes one of the:

reactants or products.

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Copper as a Reactive Electrode

Copper electrodes provide an important example.

At a copper anode:

Cu → Cu²⁺ + 2e⁻

Copper atoms lose electrons.

The copper enters the solution as:

Cu²⁺ ions.

Therefore the copper anode can gradually:

lose mass.

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Copper at the Cathode

Copper ions can move toward the cathode.

At the cathode:

Cu²⁺ + 2e⁻ → Cu

Copper ions gain electrons and become copper atoms.

Copper is deposited on the cathode.

The cathode therefore:

gains mass.

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Electrolysis with Copper Electrodes

Consider copper electrodes placed in a solution containing Cu²⁺ ions.

At the anode:

Cu → Cu²⁺ + 2e⁻

At the cathode:

Cu²⁺ + 2e⁻ → Cu

Copper is effectively transferred from:

the anode to the cathode.

This principle is important in:

• copper purification
• electroplating
• industrial electrochemistry

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Electrode Mass Changes

Electrode reactions can sometimes be detected by measuring:

mass changes.

If metal atoms leave an electrode:

electrode mass decreases.

If metal ions are deposited onto an electrode:

electrode mass increases.

For example, in a copper transfer system:

Anode → loses copper → mass decreases

Cathode → gains copper → mass increases

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Inert vs Reactive Electrodes

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Why Electrode Material Matters

Suppose the same electrolyte is electrolyzed using:

different electrode materials.

The products may not always be the same.

With inert electrodes, reactions generally involve species from the:

electrolyte.

With reactive electrodes, the electrode itself may participate.

Therefore, when analyzing electrolysis, always ask:

What is the electrode made from?

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Electrodes in Electroplating

Electroplating uses electrolysis to coat an object with a thin layer of material, usually a:

metal.

The object being coated is normally the:

cathode.

Metal ions gain electrons at its surface:

Mⁿ⁺ + ne⁻ → M

The metal forms a coating on the object.

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Example: Copper Electroplating

Suppose an object is being coated with copper.

The object is connected as the:

cathode.

At its surface:

Cu²⁺ + 2e⁻ → Cu

Copper atoms form a layer on the object.

A copper anode may also be used.

At the copper anode:

Cu → Cu²⁺ + 2e⁻

This helps replace Cu²⁺ ions removed from the solution.

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Electrodes in Metal Purification

Reactive electrodes are important in:

electrorefining.

Copper purification is a common example.

The impure copper is used as the:

anode.

A sheet of pure copper is used as the:

cathode.

Copper is transferred electrochemically from the impure anode toward the pure cathode.

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The Impure Copper Anode

At the anode:

Cu → Cu²⁺ + 2e⁻

Copper atoms become ions and enter the:

electrolyte.

The anode gradually becomes:

smaller.

Some impurities behave differently and may collect below the anode or remain in solution, depending on their chemical properties.

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The Pure Copper Cathode

At the cathode:

Cu²⁺ + 2e⁻ → Cu

Copper ions are deposited as:

solid copper.

The cathode therefore becomes:

larger and heavier.

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Analyzing an Electrode Reaction

When given an electrolysis problem, use a systematic method.

Step 1: Identify the ions present.

Step 2: Determine whether each ion is positive or negative.

Step 3: Move cations toward the cathode.

Step 4: Move anions toward the anode.

Step 5: Remember reduction occurs at the cathode.

Step 6: Remember oxidation occurs at the anode.

Step 7: Determine whether the electrodes are inert or reactive.

Step 8: Write balanced half-equations.

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Balancing Cathode Half-Equations

Suppose:

Al³⁺ → Al

The aluminum ion must gain three electrons to become neutral.

Therefore:

Al³⁺ + 3e⁻ → Al

Check the charge:

Left side:

+3 + (−3) = 0

Right side:

0

The equation is balanced.

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Balancing Anode Half-Equations

Suppose chloride ions form chlorine gas.

Start with:

Cl⁻ → Cl₂

Because chlorine gas contains two chlorine atoms:

2Cl⁻ → Cl₂

Now balance the charge.

Left side has charge:

−2

Therefore two electrons must appear on the right:

2Cl⁻ → Cl₂ + 2e⁻

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Worked Example 1

An electrolyte contains Mg²⁺ ions.

Which electrode will they move toward?

Mg²⁺ is:

positive.

Therefore:

Mg²⁺ → negative cathode.

At the cathode:

Mg²⁺ + 2e⁻ → Mg

This is:

reduction.

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Worked Example 2

An electrolyte contains I⁻ ions.

Which electrode will they move toward?

I⁻ is:

negative.

Therefore:

I⁻ → positive anode.

At the anode:

2I⁻ → I₂ + 2e⁻

This is:

oxidation.

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Worked Example 3

Write the cathode reaction for Ag⁺.

Ag⁺ must gain one electron:

Ag⁺ + e⁻ → Ag

This is reduction because the silver ion:

gains an electron.

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Worked Example 4

Write the cathode reaction for Al³⁺.

Al³⁺ needs three electrons:

Al³⁺ + 3e⁻ → Al

The aluminum ion is:

reduced.

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Worked Example 5

Write the anode reaction for Br⁻.

Bromine exists as Br₂.

Therefore:

2Br⁻ → Br₂ + 2e⁻

Bromide ions lose electrons.

This is:

oxidation.

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Worked Example 6

A copper electrode loses mass during electrolysis.

Which electrode is it likely to be if copper atoms are forming Cu²⁺ ions?

The reaction is:

Cu → Cu²⁺ + 2e⁻

Electrons are lost.

Therefore this is oxidation.

Oxidation occurs at the:

anode.

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Worked Example 7

A metal coating appears on an electrode.

Metal ions are gaining electrons:

Mⁿ⁺ + ne⁻ → M

This is reduction.

Therefore the coating forms at the:

cathode.

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Worked Example 8

A student says:

"The positive electrode is always the anode."

Is this correct?

Not for every electrochemical cell.

In an electrolytic cell:

anode = positive

In a galvanic cell:

anode = negative

The definition that always works is:

anode = oxidation.

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Worked Example 9

A student says:

"The cathode attracts electrons because it is negative."

This is incorrect reasoning.

Electrons are supplied to the cathode through the:

external circuit.

The negative cathode attracts:

positive ions.

Those cations can then receive electrons at the electrode surface.

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Worked Example 10

Why might changing from graphite electrodes to copper electrodes change an electrolysis experiment?

Graphite is relatively inert in many conditions.

Copper can be:

reactive.

Copper atoms at the anode may oxidize:

Cu → Cu²⁺ + 2e⁻

Therefore changing electrode material can change the reactions and:

products.

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Predicting Changes at Electrodes

Observations can provide clues about electrode reactions.

You might observe:

Gas bubbles

A gaseous product is forming.

Metal coating

Metal ions are being reduced at the cathode.

Electrode gets smaller

The electrode material may be oxidized.

Electrode gets heavier

Material may be deposited onto it.

Colour changes

The concentration of coloured ions may be changing.

These observations can help identify the:

chemical reactions occurring.

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Gas Formation at Electrodes

Some electrolysis reactions produce gases.

Examples include:

• hydrogen
• oxygen
• chlorine
• bromine vapour under suitable conditions

Gas production may be observed as:

bubbles at an electrode.

The identity of the gas depends on:

• electrolyte
• ions present
• electrode material
• operating conditions

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Why Electrode Reactions Matter

Electrode reactions are central to many technologies.

They are used in:

• electroplating
• metal extraction
• metal purification
• chlorine production
• hydrogen production
• battery technology
• industrial chemical manufacturing

Understanding electrodes therefore helps explain many practical applications of:

electrochemistry.

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Common Mistake: Anode Means Positive

The anode is defined by the reaction occurring there:

oxidation.

In an electrolytic cell the anode is positive.

But in a galvanic cell the anode is:

negative.

Remember:

AnOx.

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Common Mistake: Cathode Means Negative

The cathode is defined as the electrode where:

reduction occurs.

It is negative during electrolysis but positive in a:

galvanic cell.

Remember:

RedCat.

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Common Mistake: Cations Move to the Positive Electrode

Cations are positively charged.

They are attracted toward the:

negative cathode.

Therefore:

Cations → Cathode.

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Common Mistake: Electrons Move Through the Electrolyte

Electrons move through the:

external circuit and electrodes.

Ions move through the:

electrolyte.

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Common Mistake: Electrodes Never React

Some electrodes are designed to be relatively:

inert.

Others are deliberately:

reactive.

Reactive electrodes can participate directly in the electrochemical reaction.

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Common Mistake: Graphite Is Always Completely Inert

Graphite is often treated as an inert electrode in introductory electrochemistry.

However, it can react under certain:

conditions.

"Inert" therefore means relatively unreactive in the particular electrochemical system, not chemically incapable of reacting under all circumstances.

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Common Mistake: Metal Is Always Produced at the Cathode

Reduction always occurs at the cathode.

However, the substance actually produced depends on:

• ions present
• electrolyte
• concentration
• electrode material
• other possible reactions

In aqueous solutions, for example, hydrogen may sometimes be produced instead of a:

metal.

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Check Your Understanding

1. Define an electrode.
2. Name the two electrodes in an electrolytic cell.
3. Which electrode is positive during electrolysis?
4. Which electrode is negative during electrolysis?
5. Where does oxidation occur?
6. Where does reduction occur?
7. What does AnOx mean?
8. What does RedCat mean?
9. What does OIL RIG mean?
10. What is a cation?
11. Which electrode attracts cations?
12. What is an anion?
13. Which electrode attracts anions?
14. Why do ions move toward particular electrodes?
15. Where do electrons move in an electrolysis system?
16. Where do ions move?
17. What happens to electrons at the cathode?
18. What happens to electrons at the anode reaction?
19. Write the cathode half-equation for Na⁺.
20. Write the cathode half-equation for Cu²⁺.
21. Write the cathode half-equation for Al³⁺.
22. Write the anode half-equation for Cl⁻.
23. Write the anode half-equation for Br⁻.
24. What products form during electrolysis of molten sodium chloride?
25. Describe the electrode reactions during electrolysis of molten PbBr₂.
26. What is an inert electrode?
27. Give two examples of materials commonly used as inert electrodes.
28. Why is graphite commonly used as an electrode?
29. Why is platinum useful as an electrode?
30. What is a reactive electrode?
31. How does a reactive electrode differ from an inert electrode?
32. Write the oxidation reaction for a copper anode.
33. Write the reduction reaction for Cu²⁺ at a cathode.
34. Why does a copper anode lose mass?
35. Why does a copper cathode gain mass?
36. What is electroplating?
37. Why is the object being electroplated connected as the cathode?
38. Explain how copper can be electroplated onto an object.
39. Explain how reactive copper electrodes can maintain Cu²⁺ ions in an electrolyte.
40. What is electrorefining?
41. Which electrode contains impure copper during copper purification?
42. Which electrode gains pure copper?
43. Why can changing electrode material change the products of electrolysis?
44. What might bubbles at an electrode indicate?
45. What might an increase in electrode mass indicate?
46. What might a decrease in electrode mass indicate?
47. Explain why "the anode is always positive" is incorrect.
48. Compare inert and reactive electrodes.
49. Explain how ions and electrons move through an electrolytic cell.
50. Describe a systematic method for predicting the reactions at the electrodes during electrolysis.

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Key Terms

Electrode: Electrical conductor where electron-transfer reactions occur in an electrochemical system.

Anode: Electrode where oxidation occurs.

Cathode: Electrode where reduction occurs.

Cation: Positively charged ion that moves toward the cathode during electrolysis.

Anion: Negatively charged ion that moves toward the anode during electrolysis.

Oxidation: Loss of electrons.

Reduction: Gain of electrons.

Half-equation: Equation representing oxidation or reduction and showing electrons explicitly.

Inert electrode: Electrode that provides a conducting reaction surface without normally participating significantly in the reaction.

Reactive electrode: Electrode whose material participates in the electrochemical reaction.

Graphite: Conducting form of carbon commonly used as a relatively inert electrode.

Platinum: Conductive and highly corrosion-resistant metal often used as an inert electrode.

Electroplating: Use of electrolysis to deposit a thin layer of material onto an object.

Electrorefining: Use of electrolysis to purify a metal.

Electrodeposition: Formation of a material on an electrode through an electrochemical reaction.

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Key Takeaways

• An electrolytic cell contains an anode and a cathode.
• During electrolysis, the anode is positive and the cathode is negative.
• The anode is defined as the electrode where oxidation occurs.
• The cathode is defined as the electrode where reduction occurs.
• AnOx means anode = oxidation.
• RedCat means reduction = cathode.
• Cations move toward the cathode.
• Anions move toward the anode.
• Electrons move through the external circuit while ions move through the electrolyte.
• Electrode reactions can be represented using half-equations.
• At the cathode, particles gain electrons.
• At the anode, particles lose electrons.
• Inert electrodes mainly provide a conducting surface for electron transfer.
• Graphite and platinum are commonly used as relatively inert electrodes.
• Reactive electrodes participate directly in the chemical reaction.
• A reactive metal anode may dissolve and lose mass.
• A metal deposited at the cathode can cause the cathode to gain mass.
• Electrode material can affect the products of electrolysis.
• Electrodes are central to practical processes such as electroplating, electrorefining, metal extraction, and industrial electrolysis.
• When solving electrolysis problems, identify the ions, their charges, the electrode type, and whether oxidation or reduction occurs before writing the half-equations.