Galvanic Cells (VCE SSCE Chemistry): Revision Notes

Galvanic Cells

Introduction to galvanic cells

Modern portable devices including smartphones, laptops, cameras, and hearing aids require small, portable sources of electricity. These power sources come in the form of cells and batteries, which provide the convenience of operating electrical equipment without needing to be connected to mains power.

The demand for electronic devices has driven the development of many different types of cells. These range from tiny button cells in watches and calculators to large batteries that power lighthouses. While the energy from cells and batteries may cost more than energy from other sources like fossil fuels, the convenience they provide makes them invaluable.

Galvanic cells (also called voltaic cells) are electrochemical devices that convert chemical energy directly into electrical energy. An electrochemical cell is any device where chemical energy and electrical energy can be interconverted.

Historical development of galvanic cells

Galvani's discovery

Electrochemistry originated in 1791 when Italian biologist Luigi Galvani and his assistant made an unexpected observation. While experimenting with dissected frogs, they noticed a frog's leg hanging on a copper hook would twitch when it touched an iron rail. The frog's muscles were responding to an electric shock. Galvani had discovered how to generate electrical current, though he incorrectly believed it was a type of biological "life force."

Volta's voltaic pile

Other scientists investigated whether metals were truly involved in generating electricity. After years of research, Alessandro Volta developed a device in 1800 that used chemical reactions to produce electric current.

Scientists used this technique to isolate elements including hydrogen, oxygen, sodium, potassium, calcium, boron, barium, strontium, and magnesium.

In the voltaic pile, electricity is produced through a redox reaction involving zinc metal, water, and dissolved oxygen gas. The reaction produces $\text{Zn}^{2+}$ and $\text{OH}^-$ ions.

The Daniell cell: A classic example

Structure and operation

The Daniell cell, invented by English chemist John Daniell in 1836, provides an excellent example of how galvanic cells work. This cell produces an electric current that flows through an external circuit (the wire and connected device).

When the cell operates for several hours with a light bulb connected, visible changes occur:

• The zinc metal corrodes
• The copper metal develops a dark-brown coating
• The blue copper(II) sulfate solution fades in colour

Replacing the light bulb with a galvanometer (a device that detects electric current) reveals that electrons flow from the zinc electrode through the wire to the copper electrode. Importantly, current only flows when the two halves of the cell are connected by a salt bridge, typically made from filter paper soaked in an unreactive electrolyte solution such as potassium nitrate.

Reactions in the Daniell cell

The observations from the Daniell cell reveal several key processes:

The overall reaction combines both half-equations:

$\text{Cu}^{2+}(aq) + \text{Zn}(s) \rightarrow \text{Cu}(s) + \text{Zn}^{2+}(aq)$

This is a spontaneous reaction because it occurs naturally without needing external energy to drive it. In this reaction, copper(II) ions act as the oxidising agent (or oxidant), while zinc metal acts as the reducing agent (or reductant).

The Daniell cell's historical significance

The Daniell cell was a major advance in battery technology as it provided the first reliable source of electric current. It was used almost exclusively to power early telegraph systems in England and the United States. The original design used a copper cylinder with an ox gullet (oesophagus) hanging inside containing a zinc rod. The cylinder was filled with sulfuric acid solution saturated with copper sulfate, while the ox gullet contained sulfuric acid solution. The organic membrane acted as a separator, allowing ion passage.

Energy transformations in galvanic cells

Galvanic cells versus direct reactions

When copper(II) ions react directly with zinc metal (such as when zinc is immersed in copper sulfate solution), the same overall reaction occurs:

$\text{Cu}^{2+}(aq) + \text{Zn}(s) \rightarrow \text{Cu}(s) + \text{Zn}^{2+}(aq)$

However, in this direct reaction, the chemical energy transforms directly into thermal energy, which escapes as heat into the surroundings. This is an example of a spontaneous exothermic reaction.

In contrast, a galvanic cell separates the oxidation and reduction reactions into different compartments. The electrons must travel through an external circuit, so the chemical energy converts into electrical energy instead of heat.

Components of galvanic cells

Half-cells and conjugate redox pairs

A galvanic cell consists of two half-cells, each containing an electrode in contact with a solution. In the Daniell cell, one half-cell contains $\text{Cu}(s)$ and $\text{Cu}^{2+}(aq)$, while the other contains $\text{Zn}(s)$ and $\text{Zn}^{2+}(aq)$.

The species present in each half-cell form a conjugate redox pair - an oxidising agent paired with its corresponding reduced form.

Types of electrodes

Metal electrodes: When one member of the conjugate pair is a metal, that metal typically serves as the electrode itself. For example, in the zinc half-cell, the zinc metal acts as both a reactant and the electrode.

Inert electrodes: Some redox pairs, such as $\text{Br}_2(aq)/\text{Br}^-(aq)$ and $\text{Fe}^{3+}(aq)/\text{Fe}^{2+}(aq)$, don't involve solid metals. In these cases, an unreactive electrode made of platinum or graphite is used. While platinum is more chemically inert, graphite is commonly used in educational settings due to lower cost.

Gas electrodes: When one member of the conjugate pair is a gas, a special electrode setup is required.

For example, in an $\text{H}^+(aq)/\text{H}_2(g)$ half-cell, hydrogen gas bubbles through the solution and contacts a platinum electrode.

Anode and cathode

Understanding electrode terminology is essential:

The salt bridge

The salt bridge serves a critical function in galvanic cells. It contains mobile ions that balance charges developing in the two compartments as reactions proceed.

• Cations (positive ions) in the salt bridge move toward the cathode
• Anions (negative ions) in the salt bridge move toward the anode

Without a salt bridge, one compartment would accumulate excessive negative charge while the other accumulated positive charge. This charge buildup would quickly stop the reaction. The salt bridge forms part of the internal circuit of the galvanic cell.

Writing equations for galvanic cells

Half-cell equations

When a conjugate redox pair consists of an element and its corresponding ion, writing the half-equation is straightforward. For instance, for the reduction of zinc ions:

$\text{Zn}^{2+}(aq) + 2e^- \rightarrow \text{Zn}(s)$

Half-cell equations involving polyatomic ions can be more complex. For example, the reduction of dichromate ions:

$\text{Cr}_2\text{O}_7^{2-}(aq) + 14\text{H}^+(aq) + 6e^- \rightarrow 2\text{Cr}^{3+}(aq) + 7\text{H}_2\text{O}(l)$

Overall equations

The half-equations for oxidation and reduction can be combined to give an overall equation. In balanced overall equations:

Drawing and labeling galvanic cell diagrams

Identifying components from the cell reaction

When given an overall cell reaction, you can identify and draw all cell components systematically. Consider this example:

$\text{Cu}(s) + \text{Cl}_2(g) \rightarrow \text{Cu}^{2+}(aq) + 2\text{Cl}^-(aq)$

From this equation:

• Copper metal is at the anode (copper is oxidised)
• Chlorine gas is present at the cathode (chlorine is reduced)

General rules for all galvanic cells

When drawing any galvanic cell diagram:

• Electrons flow through the external circuit from anode (negative) to cathode (positive)
• Anions flow in the internal circuit toward the anode
• Cations flow toward the cathode
• The oxidation half-equation is written at the anode
• The reduction half-equation is written at the cathode

Worked example approach

Primary cells

Types of galvanic cells

Galvanic cells fall into two categories:

Primary cells are designed to be disposable and cannot be recharged. These are often called non-rechargeable cells. They "go flat" when the cell reaction reaches equilibrium (when there's no tendency for reactants and products to change). In primary cells, the products migrate away from the electrodes or are consumed by side reactions, preventing recharging.

Secondary cells are rechargeable and designed for multiple uses. These will be covered in later study.

Applications of primary cells

The alkaline cell

The alkaline cell was developed in the late 1960s to meet growing demand for small, high-capacity power sources for portable devices.

Structure: The alkaline cell is similar to simple laboratory galvanic cells but designed so both half-reactions occur in separate locations within one container. A potassium hydroxide electrolyte performs the same function as a salt bridge in simple cells.

Performance: Alkaline cells typically produce about 1.5 V. They are especially cost-effective for applications requiring high currents intermittently, such as torches and motorised toys.

Limitations: Once the reactants are consumed, the cell is "flat" and cannot be reused. This is characteristic of all primary cells.

Historical application: The Overland Telegraph

Summary