Electrochemistry: Class 11 Chemistry

Key points

Electrochemistry studies the relation between chemical change and electrical energy. Many of its reactions are oxidation reduction (redox) reactions. Oxidation is the loss of electrons, or an increase in oxidation number. Reduction is the gain of electrons, or a decrease in oxidation number. The two always happen together.

The oxidation number is the charge an atom would have if all bonds were ionic. The main rules are:

• An atom in an element has oxidation number 0.
• A simple ion has an oxidation number equal to its charge.
• Oxygen is usually −2 (−1 in peroxides) and hydrogen is usually +1 (−1 in metal hydrides).
• Group 1 metals are +1 and group 2 metals are +2, and fluorine is −1.
• The sum of oxidation numbers is zero in a neutral compound and equals the charge in an ion.

An oxidising agent gains electrons and is reduced, and a reducing agent loses electrons and is oxidised. In Zn + Cu²⁺ → Zn²⁺ + Cu, zinc is oxidised and is the reducing agent. Cu²⁺ is reduced and is the oxidising agent. Examples of oxidising agents are KMnO₄, K₂Cr₂O₇ and halogens. Common reducing agents are metals, H₂ and Fe²⁺.

Redox equations are balanced by the oxidation number method or the ion electron (half reaction) method. In the half reaction method, split the equation into oxidation and reduction half reactions, balance atoms other than H and O, then balance O with H₂O and H with H⁺ in acid solution, balance the charge with electrons, equalise the electrons and add the halves. For example, in acid: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O.

An electrolytic conductor allows current by movement of ions. Electrolysis is the chemical decomposition of an electrolyte by passing an electric current. The positive electrode is the anode, where oxidation occurs, and the negative electrode is the cathode, where reduction occurs in electrolysis. Cations move to the cathode and anions to the anode. In electrolysis, electrical energy produces a chemical change.

Examples of electrolysis: molten NaCl gives sodium at the cathode and chlorine at the anode (Down's cell). Concentrated brine gives hydrogen at the cathode, chlorine at the anode and NaOH in solution. Water with a little acid gives hydrogen and oxygen in a 2:1 volume ratio. Electroplating coats an object with a metal by making it the cathode in a solution of the metal ions.

Faraday's first law: the mass of a substance deposited or liberated at an electrode is proportional to the quantity of electricity passed, m = ZIt. Faraday's second law: for the same charge, masses of different substances are proportional to their chemical equivalents. One faraday (F) is the charge of one mole of electrons, about 96500 C mol⁻¹. The charge is Q = It, with I in amperes and t in seconds.

A galvanic (voltaic) cell converts chemical energy into electrical energy using a spontaneous redox reaction. In the Daniell cell, a zinc rod dips in ZnSO₄ solution and a copper rod in CuSO₄ solution. Zinc is the anode (oxidation, Zn → Zn²⁺ + 2e⁻). Copper is the cathode (reduction, Cu²⁺ + 2e⁻ → Cu). Electrons flow through the wire from zinc to copper. A salt bridge (KCl or KNO₃) completes the circuit and keeps both solutions neutral. In a galvanic cell the anode is negative and the cathode is positive.

Cell notation puts the anode on the left: Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s). A single vertical line shows a phase boundary and a double line shows the salt bridge. The cell emf is E°cell = E°cathode − E°anode, measured with a voltmeter when no current flows. The Daniell cell emf is about 1.10 V.

The standard electrode potential E° is the potential of a half cell with 1 M ions at 25 °C, measured against the standard hydrogen electrode (SHE), which is given 0.00 V. The SHE is a platinum electrode with H₂ gas at 1 atm bubbling into 1 M H⁺. Values are written as reduction potentials: Li⁺/Li −3.05 V, Na⁺/Na −2.71 V, Mg²⁺/Mg −2.37 V, Zn²⁺/Zn −0.76 V, Fe²⁺/Fe −0.44 V, H⁺/H₂ 0.00 V, Cu²⁺/Cu +0.34 V, Ag⁺/Ag +0.80 V and F₂/F⁻ +2.87 V.

The electrochemical series lists elements in order of electrode potential. A metal higher in the series (more negative E°) is a stronger reducing agent and displaces metals below it from their salt solutions. Metals above hydrogen displace H₂ from dilute acids. If E°cell is positive the reaction is spontaneous. The Nernst equation gives potential at other concentrations: E = E° − (0.0592/n) log Q at 25 °C.

Commercial cells: the dry (Leclanché) cell has a zinc case, a carbon rod and MnO₂ with NH₄Cl paste, giving 1.5 V. The lead acid battery uses Pb as anode and PbO₂ as cathode in H₂SO₄ and gives about 2 V per cell, so six cells give 12 V. It is rechargeable. A fuel cell, such as the hydrogen oxygen cell, makes electricity continuously as fuel is supplied and produces only water. Corrosion is the slow oxidation of metals by the environment. Rusting of iron needs oxygen and water, and is an electrochemical process. It is prevented by painting, galvanising with zinc, and sacrificial anodes of magnesium.

Study tip: Use the memory line 'An Ox, Red Cat': anode is oxidation, reduction is at the cathode. Draw the Daniell cell with all labels, including electron flow, ion flow in the salt bridge and the charge on each electrode. In Faraday questions go step by step: charge, moles of electrons, moles of substance, mass.

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Common mistakes

• Mixing up the anode and cathode. Remember: oxidation occurs at the anode in every cell.
• Saying the anode is positive in all cells. It is negative in a galvanic cell and positive in an electrolytic cell.
• Subtracting potentials the wrong way. E°cell = E°cathode − E°anode, and you do not multiply E° by coefficients.
• Forgetting the number of electrons per ion (z) in Faraday calculations.
• Leaving time in minutes in Q = It. Time must be in seconds.

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