MYP Chemistry Electrochemical Cell Notes

How Does a Simple Electrochemical Cell Work?

Basic Structure

A typical cell consists of:

1. Two different metal electrodes (e.g. Mg and Cu).
2. Each metal placed in a solution containing its own ions (e.g. Mg in MgSO₄(aq), Cu in CuSO₄(aq)).
3. A salt bridge connecting the two solutions.
4. A metal wire and external circuit connecting the two electrodes (possibly with a voltmeter or bulb).

Redox in the Cell – Mg/Cu Example

1. Consider a cell made from magnesium and copper: $$\mathrm{Mg}(\mathrm{~s})+\mathrm{Cu}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Mg}^{2+}(\mathrm{aq})+\mathrm{Cu}(\mathrm{~s})$$
2. This overall reaction is spontaneous (favourable), and the cell uses it to generate a current.
3. At the magnesium electrode (anode): oxidation $$\operatorname{Mg}(\mathrm{s}) \rightarrow \operatorname{Mg}^{2+}(\mathrm{aq})+2 e^{-}$$
   1. Mg loses electrons → it is oxidised.
   2. This electrode is the anode (oxidation always occurs at the anode).
   3. Released electrons enter the external wire.
4. At the copper electrode (cathode): reduction $$\mathrm{Cu}^{2+}(\mathrm{aq})+2 e^{-} \rightarrow \mathrm{Cu}(\mathrm{~s})$$
   1. Cu²⁺ gains electrons → it is reduced.
   2. This electrode is the cathode (reduction always occurs at the cathode).
   3. Solid copper is deposited on the copper electrode.
5. Electron Flow and Energy Conversion
   1. Electrons flow from anode to cathode through the external circuit (Mg → Cu).
   2. This flow of electrons provides electrical energy that can power a device (e.g. a light bulb).
   3. The driving force is the difference in tendency of the two metals to lose/gain electrons (their reduction potentials).
6. So, a voltaic cell converts the chemical energy of a spontaneous redox reaction into electrical energy (a current in the wire).

What Is the Function of the Salt Bridge?

1. The salt bridge is usually a tube or strip filled with an inert electrolyte (e.g. KNO₃ or K₂SO₄ in gel or solution).
2. It does not conduct electrons – those travel in the external wire.
3. Instead, the salt bridge:
   1. Allows ions to move between the half-cells.
   2. Maintains electrical neutrality (charge balance) in each solution.

How Do Metals and Electrolytes Affect Cell Voltage?

The Electrochemical Series and Reduction Potentials

1. Not all metals are equally willing to lose electrons.
2. This tendency is described by standard reduction potentials, E°, typically listed in an electrochemical series.
   1. Reactions are written as reductions (gain of electrons).
   2. A more positive E° value → species is more easily reduced (better oxidising agent).
   3. A more negative E° value → species is more easily oxidised (better reducing agent).
3. When two half-cells are combined into a cell:
   1. The half-cell with the higher E° becomes the cathode (reduction).
   2. The half-cell with the lower E° becomes the anode (oxidation).

Calculating Standard Cell Potential $E^\ominus_\text{cell}$

1. For a cell: $$\text { anode ||cathode }$$
2. The standard cell potential is: $$E_{\text {cell }}^{\circ}=E_{\text {cathode }}^{\circ}-E_{\text {anode }}^{\circ}$$
3. If E°cell is positive, the cell reaction is spontaneous under standard conditions.

How the Electrolyte Affects the Cell

1. The electrolyte is the ionic solution in each half-cell and inside the salt bridge. Its properties affect:
   1. The concentration of ions near each electrode.
   2. The rate of the redox reactions.
   3. The actual cell voltage (which can deviate from E° if concentrations differ).
2. Important points:
   1. The electrolyte must be compatible with the half-cells (no unwanted side reactions).
   2. It should be reasonably concentrated to allow good ion movement.
   3. It should be inert – it should not itself be easily oxidised or reduced.

Electrochemical Cells in Modern Technology

Electrochemical cells are at the heart of many devices: from TV remotes and phones to electric cars and medical implants.

We broadly divide them into:

1. Primary cells – non-rechargeable.
2. Secondary cells – rechargeable.
3. Fuel cells – continuously supplied with fuel.

Primary Cells (Disposable Batteries)

1. Example: Alkaline batteries (Zn–MnO₂)
   1. Common in remote controls, toys, torches.
   2. Zinc (anode) is oxidised, manganese dioxide (cathode) is reduced.
   3. Electrolyte is an alkaline paste (often KOH).
2. Advantages:
   1. Cheap and widely available.
   2. Good shelf life.
   3. Convenient for low-power, intermittent use.
3. Limitations:
   1. Not rechargeable → become waste after use.
   2. Contain materials that must be disposed of properly to avoid pollution.

Secondary Cells (Rechargeable Batteries)

1. Examples:
   1. Lead–acid batteries – used in car starter batteries.
   2. Nickel–cadmium (NiCd) – older power tools (now less common due to toxicity issues).
   3. Lithium-ion batteries – phones, laptops, electric vehicles.
2. Advantages:
   1. Rechargeable – can be used many times, reducing waste.
   2. Can deliver high current for power tools or car starters.
   3. Li-ion has high energy density → lots of energy for its mass (ideal for portable electronics and EVs).
3. Limitations:
   1. Limited lifetime (number of charge–discharge cycles).
   2. Some use toxic or rare metals (e.g. cadmium, cobalt).
   3. Require careful charging control to avoid overheating or damage.
   4. Recycling and disposal must be handled properly to avoid environmental harm.

Fuel Cells (e.g. Hydrogen Fuel Cells)

1. Fuel cells use a continuous supply of fuel (like hydrogen) and an oxidant (usually oxygen from air) to produce:
   1. Electricity
   2. Water (often the main by-product)
2. Advantages:
   1. High efficiency compared to combustion engines.
   2. Low emissions – hydrogen fuel cells produce mainly water vapour.
   3. Useful for vehicles, backup power systems, and portable power.
3. Limitations:
   1. Hydrogen production often requires a lot of energy (e.g. electrolysis), which may not be from renewable sources.
   2. Requires infrastructure for safe storage and distribution of hydrogen.
   3. Fuel cell systems can be expensive and complex.