Core practicals 10-12: electrochemical cells, redox titration and complexes

Welcome to the Lab: Mastering Core Practicals 10-12

In this chapter, we are bridging the gap between the theory of redox and transition metals and the actual hands-on work in the lab. These three practicals are frequently assessed in Paper 3, which tests your "conceptual and theoretical understanding of experimental methods."

Whether you love the lab or find it a bit daunting, these notes will break down exactly what you need to know, from how to set up a salt bridge to why your titration might turn a mysterious shade of brown. Don't worry if these seem complex at first; we will take them one step at a time!

Note: For basic titration techniques and molarity calculations, see "Core practicals 1-3". For preparation of organic substances, see "Core practicals 15-16".

Core Practical 10: Investigating Electrochemical Cells

This practical is all about building "batteries" and measuring their potential. We compare different metal/metal-ion systems to see how easily they lose or gain electrons.

1. The Setup

To measure a standard electrode potential, we need two half-cells connected together. There are three main types you might encounter:

1. Metal/Metal-ion: A strip of metal (the electrode) sitting in a solution of its own ions, e.g., \(Cu(s)\) in \(Cu^{2+}(aq)\).
2. Non-metal/Non-metal-ion: Usually involves a gas being bubbled over a platinum electrode, e.g., \(Cl_2(g)\) in \(Cl^-(aq)\).
3. Ion/Ion: A solution containing the same element in two different oxidation states, e.g., \(Fe^{2+}(aq)\) and \(Fe^{3+}(aq)\). This also requires a platinum electrode because there is no solid metal to act as a surface.

2. The Salt Bridge: The "Liquid Wire"

The salt bridge is a vital piece of equipment. It is usually a piece of filter paper soaked in an unreactive electrolyte like \(KNO_3(aq)\).

Why do we need it?
- It completes the circuit.
- It allows ions to flow between the half-cells to balance the charge.
- Crucial: We use \(KNO_3\) because potassium and nitrate ions are very unreactive and won't form precipitates with the ions in the half-cells.

3. Standard Conditions

To get the "Standard" Electrode Potential (\(E^{\theta}\)), we must maintain:

- Temperature: \(298 \text{ K}\) (\(25^{\circ}\text{C}\)).
- Pressure: \(100 \text{ kPa}\) (for gases).
- Concentration: \(1.00 \text{ mol dm}^{-3}\) for all ions in solution.

4. Measuring the Potential

We connect the two electrodes to a high-resistance voltmeter. The "high resistance" is important because it ensures that no current actually flows, allowing us to measure the maximum possible potential difference (the EMF).

Quick Review: Cell Diagrams
Remember the convention: the Standard Hydrogen Electrode (SHE) always goes on the left if used. In a general cell, the more negative \(E^{\theta}\) half-cell is usually placed on the left. The double vertical lines \(||\) represent the salt bridge.

Key Takeaway: Electrochemical cells convert chemical energy into electrical energy. The salt bridge and standard conditions are the most likely "practical" questions you'll face.

Core Practical 11: Redox Titration

Redox titrations are just like acid-base titrations, but instead of a proton transfer, we are tracking electron transfer. The two main versions you need to know involve Manganate(VII) and Iodine/Thiosulfate.

1. Potassium Manganate(VII) and Iron(II)

In this reaction, \(MnO_4^-\) is reduced to \(Mn^{2+}\), and \(Fe^{2+}\) is oxidized to \(Fe^{3+}\).

\(MnO_4^-(aq) + 8H^+(aq) + 5Fe^{2+}(aq) \rightarrow Mn^{2+}(aq) + 4H_2O(l) + 5Fe^{3+}(aq)\)

The Procedure:
- The \(MnO_4^-\) (purple) goes in the burette.
- The \(Fe^{2+}\) (unknown) goes in the conical flask.
- Self-indicating: You don't need a separate indicator! The first permanent pale pink color in the flask marks the end point.
- Acidification: You must add an excess of dilute sulfuric acid (\(H_2SO_4\)) to the flask. If you don't add enough acid, a brown precipitate of \(MnO_2\) forms, which ruins the titration.

2. Iodine and Thiosulfate

This is often used to find the concentration of an oxidizing agent (like \(Cu^{2+}\) or \(ClO^-\)).

Step 1: Add excess \(I^-\) to the oxidizing agent to produce \(I_2\). This turns the solution brown.
Step 2: Titrate the \(I_2\) with sodium thiosulfate (\(Na_2S_2O_3\)).
\(I_2(aq) + 2S_2O_3^{2-}(aq) \rightarrow 2I^-(aq) + S_4O_6^{2-}(aq)\)

The Indicator Trick:
- As you titrate, the brown iodine color fades to straw-yellow.
- When it is straw-yellow, add starch indicator. The solution turns blue-black.
- Continue adding thiosulfate dropwise until the blue-black color disappears (colorless). Adding starch too early can trap iodine and give an inaccurate result!

3. Common Calculations

1. Find the moles of the "known" solution (\(n = c \times V\)).
2. Use the balanced equation ratio to find moles of the "unknown".
3. Scale up (e.g., if you only titrated \(25 \text{ cm}^3\) out of a \(250 \text{ cm}^3\) volumetric flask).
4. Calculate concentration or percentage purity.

Key Takeaway: Always watch for the color change. For Manganate, it's colorless to pink. For Iodine/Starch, it's blue-black to colorless.

Core Practical 12: Preparation of a Transition Metal Complex

Transition metals are famous for forming complex ions with ligands. In this practical, you usually synthesize a solid salt of a complex, such as tetraamminecopper(II) sulfate 1-hydrate.

1. The Chemistry

Most preparations involve ligand exchange. For example, replacing water ligands around a copper ion with ammonia ligands:

\([Cu(H_2O)_6]^{2+}(aq) + 4NH_3(aq) \rightarrow [Cu(NH_3)_4(H_2O)_2]^{2+}(aq) + 4H_2O(l)\)

2. The Practical Steps

1. Dissolving: Dissolve the metal salt in a minimum amount of solvent.
2. Ligand Addition: Add the ligand (e.g., concentrated ammonia) in a fume cupboard.
3. Precipitation: Add a "cold" solvent like ethanol. Complexes are often less soluble in alcohol, which forces the solid to "crash out" of the solution.
4. Filtration: Use vacuum filtration (a Buchner funnel and flask). This is much faster than gravity filtration and leaves the solid much drier.
5. Washing: Wash the crystals with a little cold ethanol to remove impurities without dissolving the product.
6. Drying: Leave the crystals in a warm place or a desiccator.

3. Evaluating Success

You will often be asked to calculate the percentage yield:

\(\text{Percentage Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%\)

Common reasons for a low yield:
- Loss of product during filtration.
- The reaction not going to completion.
- Product remaining dissolved in the solvent.

Key Takeaway: Focus on the purification steps—vacuum filtration and washing with cold solvent are high-frequency exam topics.

Quick Summary Table

CP 10 (Cells): Measure \(E_{cell}\). Needs Salt Bridge (\(KNO_3\)) and Standard Conditions.
CP 11 (Redox): \(MnO_4^-\) (pink end point) or \(I_2\)/Thiosulfate (starch indicator).
CP 12 (Complexes): Synthesis by ligand exchange. Uses Buchner filtration and ethanol precipitation.

Common Mistakes to Avoid

- Burette Readings: In \(MnO_4^-\) titrations, the purple color is so dark you often read the top of the meniscus rather than the bottom. Just be consistent!
- Standard Units: Ensure concentrations are in \(1.00 \text{ mol dm}^{-3}\). If a question gives you \(2.00 \text{ mol dm}^{-3}\), it is not standard.
- Calculations: In the \(I_2/S_2O_3^{2-}\) titration, remember the 1:2 ratio. One mole of \(I_2\) reacts with two moles of thiosulfate.
- Safety: Transition metal salts and concentrated ammonia are often toxic or corrosive. Always mention gloves and fume cupboards in "Risk and Hazard" questions.