What Is a Half Cell?

A half cell is one part of an electrochemical cell.

A typical metal-ion half cell contains:

• an electrode
• an electrolyte
• a particular oxidation or reduction half-reaction

For example, a zinc half cell can contain:

a zinc electrode in a solution containing Zn²⁺ ions

A copper half cell can contain:

a copper electrode in a solution containing Cu²⁺ ions

Two suitable half cells can be connected to form a complete:

electrochemical cell.

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Why Is It Called a Half Cell?

A complete redox reaction contains two processes:

oxidation

and:

reduction.

A half cell represents one side of this overall electron-transfer process.

When two half cells are connected:

one half cell undergoes oxidation

while:

the other undergoes reduction.

Together they form the complete redox reaction.

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The Two Main Components

For a simple metal/metal-ion half cell, the two main physical components are:

1. Electrode

2. Electrolyte

Each has a different function.

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The Electrode

An electrode is an electrical conductor where oxidation or reduction occurs.

Electrodes are often made from:

metals.

Examples include:

• zinc
• copper
• silver
• iron

The electrode provides a surface where electron transfer can occur.

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The Electrolyte

An electrolyte is a substance containing mobile:

ions.

In a half cell, the electrolyte is commonly an aqueous solution containing ions related to the electrode.

For example:

Zinc half cell:

Zn electrode + Zn²⁺ solution

Copper half cell:

Cu electrode + Cu²⁺ solution

Silver half cell:

Ag electrode + Ag⁺ solution

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A Zinc Half Cell

A simple zinc half cell contains:

Zn(s) electrode

placed in a solution containing:

Zn²⁺(aq).

We can represent the system as:

Zn(s) | Zn²⁺(aq)

The vertical line represents the boundary between two:

phases.

In this case:

solid zinc | aqueous zinc ions

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What Can Happen in the Zinc Half Cell?

Zinc atoms can lose electrons:

Zn(s) → Zn²⁺(aq) + 2e⁻

This is:

oxidation.

Remember:

Oxidation = loss of electrons

As zinc atoms become Zn²⁺ ions, they leave the electrode and enter the:

solution.

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A Copper Half Cell

A simple copper half cell contains:

Cu(s) electrode

in a solution containing:

Cu²⁺(aq).

It can be represented as:

Cu²⁺(aq) | Cu(s)

when written as the reduction half-cell on the right side of a conventional cell diagram.

Copper ions can gain electrons:

Cu²⁺(aq) + 2e⁻ → Cu(s)

This is:

reduction.

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Oxidation and Reduction in Half Cells

A half-reaction can be written in either direction.

For example:

Zn → Zn²⁺ + 2e⁻

shows zinc being:

oxidized.

But:

Zn²⁺ + 2e⁻ → Zn

shows zinc ions being:

reduced.

Which direction actually occurs depends on what other half cell it is connected to and the conditions of the system.

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Half-Reactions

A half-reaction shows either oxidation or reduction separately.

For example:

Zn(s) → Zn²⁺(aq) + 2e⁻

is an oxidation half-reaction.

Cu²⁺(aq) + 2e⁻ → Cu(s)

is a reduction half-reaction.

The electrons appear explicitly in half-reactions.

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Connecting Two Half Cells

Suppose we connect:

Zn/Zn²⁺

and:

Cu²⁺/Cu.

Zinc has the greater tendency to be oxidized under the usual cell conditions.

Therefore:

Zn → Zn²⁺ + 2e⁻

Copper ions accept the electrons:

Cu²⁺ + 2e⁻ → Cu

The two half-reactions work together.

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Following the Electron Transfer

The overall zinc-copper reaction is:

Zn + Cu²⁺ → Zn²⁺ + Cu

The zinc loses electrons while the copper ions gain them.

The electron transfer underlying this reaction can be visualized here:

This demonstrates an important principle:

oxidation and reduction always occur together.

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Why Must Two Half-Reactions Occur?

Electrons cannot simply appear or disappear during a chemical reaction.

If one substance:

loses electrons,

another substance must:

gain those electrons.

Therefore:

oxidation requires reduction

and:

reduction requires oxidation.

This is why the overall process is called a:

redox reaction.

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Anode and Cathode

When two half cells form a galvanic cell, the electrodes are called the:

anode

and:

cathode.

The definitions are:

Anode = oxidation

Cathode = reduction

A useful memory aid is:

AN OX — ANode OXidation

RED CAT — REDuction CAThode

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The Anode Half Cell

At the anode, atoms or other species:

lose electrons.

For the zinc-copper cell:

Zn → Zn²⁺ + 2e⁻

Therefore, the zinc half cell acts as the:

anode half cell.

In a galvanic cell, the anode is:

negative.

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The Cathode Half Cell

At the cathode, species:

gain electrons.

For the zinc-copper cell:

Cu²⁺ + 2e⁻ → Cu

Therefore, the copper half cell acts as the:

cathode half cell.

In a galvanic cell, the cathode is:

positive.

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Electron Flow Between Half Cells

When the two half cells are connected through an external circuit, electrons travel:

from the anode to the cathode.

For a zinc-copper cell:

Zn electrode → wire → Cu electrode

The electrons released by zinc are eventually accepted by:

Cu²⁺ ions.

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What Happens to the Zinc Electrode?

At the zinc electrode:

Zn → Zn²⁺ + 2e⁻

Solid zinc becomes aqueous zinc ions.

Therefore:

• zinc atoms leave the electrode
• Zn²⁺ concentration increases
• the zinc electrode loses mass

The electrode gradually becomes:

smaller.

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What Happens to the Copper Electrode?

At the copper electrode:

Cu²⁺ + 2e⁻ → Cu

Copper ions become solid copper atoms.

Therefore:

• Cu²⁺ ions leave the solution
• Cu²⁺ concentration decreases
• copper is deposited on the electrode
• the electrode gains mass

The electrode gradually becomes:

larger.

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Why Do We Need a Salt Bridge?

Connecting the electrodes with a wire is not enough.

As the reactions continue, electrical charge would build up in the half-cell solutions.

A:

salt bridge

allows ions to move between the half cells.

This helps maintain electrical neutrality and allows the redox reaction to continue.

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The Salt Bridge Does Not Carry Electrons

This distinction is important.

External wire → electrons

Salt bridge → ions

Electrons travel through the metal wire.

Ions move through the electrolytes and salt bridge.

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Different Types of Half Cells

Not every half cell consists of a metal placed in a solution containing its ions.

Several types are possible.

These include:

• metal/metal-ion half cells
• gas half cells
• ion/ion half cells
• half cells involving insoluble salts

Sometimes an inert electrode is needed.

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Metal/Metal-Ion Half Cells

These are among the simplest half cells.

Examples include:

Zn(s) | Zn²⁺(aq)

Cu(s) | Cu²⁺(aq)

Ag(s) | Ag⁺(aq)

Fe(s) | Fe²⁺(aq)

The metal itself acts as the:

electrode.

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Silver Half Cell

A silver half cell may contain:

Ag(s)

and:

Ag⁺(aq).

The reduction half-reaction is:

Ag⁺(aq) + e⁻ → Ag(s)

The oxidation reaction would be the reverse:

Ag(s) → Ag⁺(aq) + e⁻

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Iron Half Cell

An iron half cell may contain:

Fe(s)

and:

Fe²⁺(aq).

The reduction half-reaction can be written:

Fe²⁺(aq) + 2e⁻ → Fe(s)

Again, whether this actually occurs as reduction depends on the other half cell to which it is connected.

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Comparing Metal Half Cells

Consider:

Zn/Zn²⁺

Cu/Cu²⁺

Ag/Ag⁺

Each contains:

• a metal electrode
• ions of that metal in solution
• an oxidation/reduction pair

However, the metals have different tendencies to:

lose or gain electrons.

This difference allows different combinations of half cells to produce different cell voltages.

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Half Cells Without a Solid Reactive Metal

Sometimes both members of a redox pair are dissolved:

ions.

For example:

Fe³⁺ and Fe²⁺.

The reaction is:

Fe³⁺ + e⁻ ⇌ Fe²⁺

Neither Fe³⁺ nor Fe²⁺ provides a solid conducting electrode.

Therefore, an:

inert electrode

is required.

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Inert Electrodes

An inert electrode conducts electrons but does not normally take part directly in the overall redox reaction.

Common materials include:

• platinum
• graphite

For an Fe³⁺/Fe²⁺ half cell, a platinum electrode could provide the conducting surface.

The electrode transfers electrons while the iron ions undergo the chemical change.

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Gas Half Cells

Some half cells involve a:

gas.

One important example is the hydrogen half cell.

The half-reaction is:

2H⁺(aq) + 2e⁻ ⇌ H₂(g)

Because neither H⁺ ions nor hydrogen gas provides a suitable solid electrode, an inert:

platinum electrode

is commonly used.

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The Hydrogen Half Cell

A hydrogen half cell contains:

• hydrogen gas
• H⁺ ions in solution
• a platinum electrode

The platinum provides a surface for electron transfer.

It does not supply the hydrogen atoms involved in the:

redox reaction.

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Standard Hydrogen Electrode

The standard hydrogen electrode, or SHE, is an important reference half cell.

Its standard electrode potential is defined as:

0.00 V.

Other electrode potentials can be measured relative to this reference.

This allows scientists to compare the tendencies of different half cells to undergo:

reduction.

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Half-Cell Potential

An isolated half cell does not provide a directly measurable voltage by itself.

Voltage is measured as a:

difference in potential between two half cells.

Therefore, electrode potentials are measured relative to a reference electrode.

This is why the standard hydrogen electrode is important.

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Standard Conditions

Standard electrode potentials are normally defined under specified standard conditions.

Commonly these include:

• dissolved species at standard-state activity, often approximated in school chemistry as 1 mol dm⁻³
• gases at standard pressure
• a specified temperature, commonly 25°C for tabulated values

Keeping conditions controlled allows meaningful comparisons between different:

half cells.

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Simple Half-Cell Diagrams

A half cell can be represented using:

cell notation.

For example:

Zn(s) | Zn²⁺(aq)

The symbols tell us:

Zn(s) = solid zinc electrode

| = phase boundary

Zn²⁺(aq) = zinc ions dissolved in water

This is a compact way of describing the half cell.

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Understanding State Symbols

Cell notation often uses state symbols.

(s) = solid

(l) = liquid

(g) = gas

(aq) = dissolved in water

For example:

Cu(s) | Cu²⁺(aq)

contains solid copper and aqueous copper ions.

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Phase Boundaries

A single vertical line:

|

represents a boundary between different:

phases.

For example:

Zn(s) | Zn²⁺(aq)

means:

solid zinc in contact with aqueous zinc ions.

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Complete Cell Diagrams

Two half cells can be combined into a complete cell diagram.

For the zinc-copper cell:

Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

The double vertical line:

||

represents the connection between the electrolyte solutions, commonly the:

salt bridge.

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Reading a Cell Diagram

Consider:

Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

Reading from left to right:

Zn(s) — zinc electrode

| — phase boundary

Zn²⁺(aq) — zinc ions

|| — salt bridge

Cu²⁺(aq) — copper ions

| — phase boundary

Cu(s) — copper electrode

This represents two connected:

half cells.

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Conventional Cell Notation

For a spontaneous galvanic cell, the conventional arrangement is generally:

anode | anode solution || cathode solution | cathode

Therefore:

oxidation is shown on the left

and:

reduction is shown on the right.

For the zinc-copper cell:

Zn | Zn²⁺ || Cu²⁺ | Cu

Zinc is oxidized and Cu²⁺ is reduced.

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Constructing a Cell Diagram

Suppose we have:

Mg/Mg²⁺

and:

Cu²⁺/Cu.

If magnesium is the anode and copper is the cathode, the diagram is:

Mg(s) | Mg²⁺(aq) || Cu²⁺(aq) | Cu(s)

This immediately communicates the arrangement of the electrochemical cell.

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Interpreting a Cell Diagram

Consider:

Fe(s) | Fe²⁺(aq) || Ag⁺(aq) | Ag(s)

If written as a spontaneous galvanic cell:

Left side:

oxidation

Fe → Fe²⁺ + 2e⁻

Right side:

reduction

Ag⁺ + e⁻ → Ag

The half-reactions must then be balanced for electrons before forming the overall equation.

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Balancing the Electrons

For:

Fe → Fe²⁺ + 2e⁻

and:

Ag⁺ + e⁻ → Ag

the silver half-reaction must occur twice:

2Ag⁺ + 2e⁻ → 2Ag

Now the electrons cancel.

Overall:

Fe + 2Ag⁺ → Fe²⁺ + 2Ag

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Comparing Zinc and Copper Half Cells

Zinc half cell

Electrode:

Zn

Electrolyte contains:

Zn²⁺

When acting as the anode:

Zn → Zn²⁺ + 2e⁻

Electrode mass:

decreases

Copper half cell

Electrode:

Cu

Electrolyte contains:

Cu²⁺

When acting as the cathode:

Cu²⁺ + 2e⁻ → Cu

Electrode mass:

increases

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Comparing Copper and Silver Half Cells

Copper reduction:

Cu²⁺ + 2e⁻ → Cu

Silver reduction:

Ag⁺ + e⁻ → Ag

Both involve:

metal ions gaining electrons and becoming solid metal.

However, the ions require different numbers of electrons.

Cu²⁺ requires:

2 electrons.

Ag⁺ requires:

1 electron.

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Comparing Metal and Gas Half Cells

Metal/metal-ion half cell

Example:

Cu²⁺/Cu

The metal itself can act as the electrode.

Gas half cell

Example:

H⁺/H₂

There is no solid hydrogen electrode.

An inert conductor such as platinum is therefore required.

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Comparing Reactive and Inert Electrodes

A reactive electrode participates directly in the half-reaction.

Example:

Zn → Zn²⁺ + 2e⁻

The zinc electrode is consumed.

An inert electrode provides a conducting surface without being consumed in the main redox reaction.

Example:

Pt electrode in an Fe³⁺/Fe²⁺ system.

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Dynamic Processes at an Electrode

At the boundary between an electrode and electrolyte, particles can undergo electron-transfer reactions.

For a metal-ion system:

Mⁿ⁺ + ne⁻ ⇌ M

The double arrow reminds us that oxidation or reduction may be possible depending on:

• the other half cell
• concentrations
• temperature
• electrode potentials

The direction of the overall reaction depends on the complete system.

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Concentration Matters

The behavior and potential of a half cell can be affected by the concentration of ions in the:

electrolyte.

For example, changing the concentration of Cu²⁺ can change the potential of the:

Cu²⁺/Cu half cell.

This becomes important when studying:

cell potentials and the Nernst equation.

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Temperature Matters

Electrode potentials can also depend on:

temperature.

For accurate comparisons, experimental conditions should therefore be controlled.

This is why standard electrode potentials are defined under specified:

standard conditions.

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Half Cells in Batteries

Real batteries contain systems that function according to the same fundamental principles.

Different materials are used to create:

• oxidation reactions
• reduction reactions
• electron flow
• ion movement

The chemistry may be more complex than the simple Zn/Cu system, but the core idea remains:

two coupled half-reactions produce an electrochemical process.

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Worked Example 1

Identify the electrode and electrolyte in:

Zn(s) | Zn²⁺(aq)

Electrode:

Zn(s)

Electrolyte:

solution containing Zn²⁺ ions

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Worked Example 2

Write the oxidation half-reaction for zinc.

Zn(s) → Zn²⁺(aq) + 2e⁻

Zinc:

loses electrons.

Therefore, zinc is:

oxidized.

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Worked Example 3

Write the reduction half-reaction for silver ions.

Ag⁺(aq) + e⁻ → Ag(s)

Silver ions:

gain electrons.

Therefore, Ag⁺ is:

reduced.

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Worked Example 4

Interpret:

Cu(s) | Cu²⁺(aq)

This represents:

a copper electrode in contact with a solution containing Cu²⁺ ions.

The vertical line represents the:

phase boundary.

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Worked Example 5

Interpret:

Zn | Zn²⁺ || Cu²⁺ | Cu

This represents:

a zinc half cell connected to a copper half cell.

The double line represents the:

salt bridge or liquid junction.

Under the conventional spontaneous-cell notation:

Zinc undergoes:

oxidation.

Cu²⁺ undergoes:

reduction.

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Worked Example 6

A metal electrode becomes lighter as a galvanic cell operates.

What may be happening?

Metal atoms are probably being:

oxidized into ions.

For example:

M → Mⁿ⁺ + ne⁻

The atoms leave the electrode and enter the solution.

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Worked Example 7

A metal electrode becomes heavier.

What may be happening?

Metal ions are probably being:

reduced and deposited on the electrode.

For example:

Mⁿ⁺ + ne⁻ → M

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Worked Example 8

A half cell contains Fe³⁺ and Fe²⁺ ions but no solid iron.

Why might platinum be required?

The system needs a conducting surface for:

electron transfer.

Platinum can act as an:

inert electrode.

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Common Mistake: A Half Cell Is Half of a Container

The term half cell does not simply mean half of the physical apparatus.

It refers to one part of the:

overall redox system.

Each half cell is associated with one oxidation/reduction couple.

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Common Mistake: Every Half Cell Uses a Metal Electrode

Many simple half cells use metal electrodes.

However, some redox systems do not contain a suitable solid conductor.

These require an:

inert electrode.

Examples include platinum or graphite.

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Common Mistake: Oxidation Always Happens in a Particular Half Cell

A redox pair can potentially operate in either direction.

For example:

Cu → Cu²⁺ + 2e⁻

or:

Cu²⁺ + 2e⁻ → Cu

Which occurs depends on the other half cell and the conditions.

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Common Mistake: Electrons Move Through the Electrolyte

Electrons mainly move through the:

external metallic circuit.

Charge within the electrolyte is carried by:

ions.

This distinction is fundamental to electrochemistry.

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Common Mistake: The Salt Bridge Supplies Electrons

The salt bridge supplies a pathway for:

ions.

It does not provide electrons to the electrodes.

Electrons travel through the:

external circuit.

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Common Mistake: The Electrode and Electrolyte Are the Same Thing

The:

electrode

is the conducting solid surface.

The:

electrolyte

contains mobile ions.

They work together but have different functions.

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Common Mistake: A Single Half Cell Has a Measurable Absolute Voltage

Electrode potential is measured relative to another:

half cell.

There is no experimentally measurable absolute electrode potential for an isolated half cell in ordinary electrochemistry.

This is why reference electrodes are used.

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Check Your Understanding

1. Define a half cell.

2. Name the two main components of a simple metal-ion half cell.

3. What is an electrode?

4. What is an electrolyte?

5. What particles carry charge through an electrolyte?

6. Describe a Zn/Zn²⁺ half cell.

7. Describe a Cu²⁺/Cu half cell.

8. Write the oxidation half-reaction for zinc.

9. Write the reduction half-reaction for Cu²⁺.

10. Define oxidation.

11. Define reduction.

12. Why must oxidation and reduction occur together?

13. What is a half-reaction?

14. Where does oxidation occur in a galvanic cell?

15. Where does reduction occur?

16. What is the anode?

17. What is the cathode?

18. In which direction do electrons flow through a galvanic cell?

19. Why does a zinc anode lose mass?

20. Why can a copper cathode gain mass?

21. What is the purpose of a salt bridge?

22. What particles move through a salt bridge?

23. What does Zn(s) | Zn²⁺(aq) represent?

24. What does the single vertical line in cell notation mean?

25. What does the double vertical line usually represent?

26. Interpret Zn | Zn²⁺ || Cu²⁺ | Cu.

27. Why is oxidation normally written on the left in conventional galvanic-cell notation?

28. Write a cell diagram for a magnesium-copper galvanic cell in which magnesium is oxidized.

29. What is an inert electrode?

30. Why might platinum be used in a half cell?

31. Describe an Fe³⁺/Fe²⁺ half cell.

32. Why does an Fe³⁺/Fe²⁺ half cell require an inert electrode?

33. Describe the basic components of a hydrogen half cell.

34. What is the standard hydrogen electrode?

35. Why are electrode potentials measured relative to a reference electrode?

36. Compare a metal/metal-ion half cell with a gas half cell.

37. Compare reactive and inert electrodes.

38. How can concentration affect a half cell?

39. A metal electrode decreases in mass. What type of process may be occurring?

40. Explain how two half cells work together to produce a complete redox reaction.

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Key Terms

• Half cell: One part of an electrochemical cell associated with an oxidation or reduction half-reaction.
• Electrode: Conducting material where electron transfer occurs.
• Electrolyte: Substance containing mobile ions capable of conducting charge.
• Half-reaction: Equation showing either oxidation or reduction separately.
• Oxidation: Loss of electrons.
• Reduction: Gain of electrons.
• Redox reaction: Reaction involving both oxidation and reduction.
• Anode: Electrode where oxidation occurs.
• Cathode: Electrode where reduction occurs.
• Metal/metal-ion half cell: Half cell containing a metal electrode in contact with ions of that metal.
• Inert electrode: Conducting electrode that provides a surface for electron transfer without participating significantly in the overall reaction.
• Platinum electrode: Common inert electrode used in some half cells.
• Salt bridge: Connection that allows ions to move between half cells and maintain electrical neutrality.
• Phase boundary: Interface between two different physical phases, represented by a single vertical line in cell notation.
• Cell notation: Compact symbolic representation of an electrochemical cell.
• Standard hydrogen electrode: Reference half cell assigned a standard electrode potential of 0.00 V.
• Electrode potential: Measure of the tendency of a half-cell redox couple to undergo reduction relative to a reference electrode.
• Aqueous: Dissolved in water.
• Galvanic cell: Electrochemical cell that produces electrical energy from a spontaneous redox reaction.

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Key Takeaways

• A half cell represents one part of an electrochemical redox system.
• A simple half cell contains an electrode and electrolyte.
• The electrode provides a conducting surface for electron transfer.
• The electrolyte contains mobile ions.
• A metal/metal-ion half cell contains a metal electrode and ions of that metal.
• Zn(s) | Zn²⁺(aq) represents a zinc half cell.
• Cu²⁺(aq) | Cu(s) can represent the copper reduction half-cell when written on the right side of a cell diagram.
• Oxidation means loss of electrons.
• Reduction means gain of electrons.
• Oxidation and reduction always occur together.
• A half-reaction shows one part of the overall redox reaction.
• The anode is where oxidation occurs.
• The cathode is where reduction occurs.
• Electrons flow through the external circuit from anode to cathode in a galvanic cell.
• An electrode undergoing oxidation may lose mass as atoms become ions.
• An electrode receiving metal ions through reduction may gain mass.
• The salt bridge allows ions, not electrons, to move between half cells.
• A single vertical line in cell notation represents a phase boundary.
• A double vertical line represents a salt bridge or liquid junction.
• A complete zinc-copper cell can be written as Zn | Zn²⁺ || Cu²⁺ | Cu.
• In conventional spontaneous galvanic-cell notation, oxidation is written on the left and reduction on the right.
• Not all half cells contain a reactive metal electrode.
• Some systems require an inert electrode, such as platinum or graphite.
• Gas half cells can also be constructed.
• The hydrogen half cell uses H⁺, H₂ gas, and an inert platinum electrode.
• The standard hydrogen electrode provides a reference for comparing electrode potentials.
• An isolated half cell does not have a directly measurable absolute electrode voltage.
• Different half-cell systems have different tendencies to gain or lose electrons.
• Combining suitable half cells allows chemical energy to be converted into electrical energy.