Page 2: Electrode Reactions and Half-Cells

This page delves deeper into the specific reactions occurring at each electrode in an electrochemical cell. It explains the concepts of oxidation and reduction half-reactions, as well as the overall cell reaction.

Definition: The anode is the electrode where oxidation occurs, while the cathode is the electrode where reduction takes place.

In the example of a zinc-copper cell:

1. Anode $Zinc electrode$: Zn$s$ → Zn²⁺$aq$ + 2e⁻ $Oxidation half-reaction$
2. Cathode $Copper electrode$: Cu²⁺$aq$ + 2e⁻ → Cu$s$ $Reduction half-reaction$
3. Overall cell reaction: Zn$s$ + Cu²⁺$aq$ → Zn²⁺$aq$ + Cu$s$

Highlight: The direction of each reaction depends on the reactivity of the metals involved, which is measured by their electrode potentials.

The page also introduces the concept of electrode potentials:

• A metal that is easily oxidized has a negative electrode potential
• A metal that is easily reduced has a positive electrode potential

Example: Half-cells can also involve solutions of two aqueous ions of the same element, such as Fe²⁺$aq$/Fe³⁺$aq$. In this case, a platinum electrode is used as an inert conductor.

Vocabulary: Inert electrode - An electrode that does not participate in the redox reaction but serves as a conductor for electron transfer.

Page 3: Understanding Electrode Potentials

This page focuses on the interpretation of electrode potentials and their relationship to the reactivity of elements in electrochemical cells. It provides crucial information for students studying electrode potentials in A-Level Chemistry.

Definition: The electrode potential is a measure of the tendency of a species to undergo oxidation or reduction.

Key points about electrode potentials:

• Negative electrode potential indicates a strong reducing agent $easily oxidized$
• Positive electrode potential indicates a strong oxidizing agent $easily reduced$

Example: In the zinc-copper cell:

• Zn²⁺/Zn has an electrode potential of -0.76 V $reducing agent, more reactive$
• Cu²⁺/Cu has an electrode potential of +0.34 V $oxidizing agent, less reactive$

The page also provides guidance on drawing electrochemical cells:

• Place the more negative $more reactive$ electrode on the left-hand side
• Show the direction of electron flow in the external circuit
• Indicate the salt bridge connecting the two half-cells

Highlight: Understanding how to draw and interpret electrochemical cell diagrams is essential for solving electrode potential A-level Chemistry questions.

Page 4: Electrochemical Series

This page presents the electrochemical series, a crucial tool for predicting the behavior of elements in electrochemical reactions. The series is essential for students studying oxidation and reduction in electrochemical cells.

Definition: The electrochemical series is a ranking of elements and ions based on their standard electrode potentials, indicating their relative tendencies to undergo oxidation or reduction.

The page provides a comprehensive table of the electrochemical series, including:

• Elements and their ions
• Electrode reactions $oxidized form + ne⁻ → reduced form$
• Standard electrode potentials $E°$ in volts

Highlight: The electrochemical series is arranged from the most negative potential $strongest reducing agent$ to the most positive potential $strongest oxidizing agent$.

Key observations from the electrochemical series:

• Lithium has the most negative potential $-3.05 V$, making it the strongest reducing agent
• Fluorine has the most positive potential $+2.87 V$, making it the strongest oxidizing agent
• Hydrogen has a standard potential of 0.00 V, serving as the reference point

Example: When comparing lithium and magnesium, lithium has a more negative potential $-3.05 V vs. -2.37 V$, indicating that it is more easily oxidized and should be placed on the left side of an electrochemical cell diagram.

Vocabulary: Standard electrode potential - The potential of a half-cell measured against the standard hydrogen electrode under standard conditions $1 M concentration, 1 atm pressure, 25°C$.

Page 5: Standard Hydrogen Electrode and Cell Notation

This page introduces the concept of the Standard Hydrogen Electrode $SHE$ and explains cell notation, both crucial topics for understanding electrode potentials in A-Level Chemistry.

Definition: The Standard Hydrogen Electrode $SHE$ is a primary reference electrode used to measure the standard electrode potentials of other half-cells.

Components of the Standard Hydrogen Electrode:

• Platinum electrode
• Hydrogen gas at 100 kPa $1 atm$ pressure
• 1.0 M H⁺$aq$ solution
• Temperature maintained at 298 K $25°C$

Highlight: The SHE is assigned a standard potential of 0.00 V, serving as the reference point for all other electrode potentials.

The page also explains cell notation, a shorthand method for representing electrochemical cells:

• Format: Reduced form | Oxidized form || Oxidized form | Reduced form
• Example: Pt$s$ | H₂$g$ | H⁺$aq$ || Cu²⁺$aq$ | Cu$s$

Vocabulary: Phase boundary - The interface between two different phases in an electrochemical cell, such as solid-liquid or gas-liquid interfaces.

Example: In a cell combining the SHE with a copper electrode: Left half-cell: Pt$s$ | H₂$g$ | H⁺$aq$ Right half-cell: Cu²⁺$aq$ | Cu$s$ Cell notation: Pt$s$ | H₂$g$ | H⁺$aq$ || Cu²⁺$aq$ | Cu$s$

Understanding the SHE and cell notation is essential for students working on setting up electrochemical cells with salt bridge lab reports and solving complex electrochemistry problems.

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Page 1: Introduction to Electrochemical Cells

This page introduces the fundamental concepts of electrochemical cells and their components. Electrochemical cells are devices that convert chemical energy into electrical energy through redox reactions. They consist of two half-cells connected by a salt bridge and an external circuit.

Definition: An electrochemical cell is a system composed of two half-cells that undergo redox reactions to generate an electric current.

The key components of an electrochemical cell include:

1. Two different metals dipped in solutions of their own ions
2. A wire connecting the metals to form an external circuit
3. A salt bridge connecting the two solutions

Highlight: The salt bridge is crucial for maintaining electrical neutrality in the cell by allowing ion flow between half-cells.

The page also explains the process of setting up an electrochemical cell:

1. Connect the two metals with a wire for electron flow
2. Join the solutions with a salt bridge for ion flow
3. Use a voltmeter in the circuit to measure the potential difference $emf$

Example: In a copper-zinc cell, zinc loses electrons more easily than copper, making it the anode where oxidation occurs. Electrons flow through the external circuit from zinc to copper, while the salt bridge allows ions to flow between half-cells to balance charges.

Vocabulary: EMF $electromotive force$ - The potential difference between the two electrodes in an electrochemical cell.