Electrochemical cells

Example 1: Nickel - Hydrogen cell with a carbon electrode

C(s) | H+(aq), Cl-(aq) || Ni2+(aq), SO42-(aq) | Ni(s)

Cathode: 2H+ + 2e- --> H2 Ešr = + 0.00V

Anode: Ni --> Ni2+ + 2e- Ešr = + 0.26V

Balanced Net Equation: 2H+ + Ni --> H2 + Ni2+ Eš = +0.26V

The reaction is spontaneous; the hydrogen ions have a stronger pull on the electrons in nickel than nickel atoms themselves have, resulting in a reaction. The voltage detected by the voltmeter confirms there is a spontaneous reaction, as well as the presence of bubbles at the cathode.

Direction of electron flow: anode (-ve) --> cathode (+ve)

 

Descriptions of reactants and products:

HCl - liquid, colourless, clear, acidic odour

C - solid, black, dull, brittle

NiSO4 - liquid, green, clear, odourless

Ni - solid, bronze colour, lustrous, malleable

Product - Bubbles appear around the carbon electrode. A positive voltage was detected.

Identification of product: From the net equation, the products are H2 and Ni2+. Bubbles have been oberserved at the cathode, indicating the presence of H2(g). Ni2+ is present in the NiSO4 solution, indicating the presence of the ion.

Nickel-Hydrogen cell
Full Cell
Welcome to the electrochemical cells page. Electrochemical cells are a special application of redox reactions. Essentially, two 'containers' are set up, each with a separate electrolyte (a solution containing ions); the electrodes (a solid conductor, either a non-reactive metal or carbon, or the solid form of the electrolyte) are connected by a wire. What happens then is based on the difference in reduction potentials: in an electrochemical cell, also called a galvanic cell, the reduction potential is positive. This means that one of the 'containers' has a SOA, which undergoes a reduction reaction, and has a stronger pull on the electrons in the other container than the SRA in the other container has; the electrons are drawn through the wire, creating a current. The electrode which gains electrons and reduces the electrolyte in its cell is called the cathode, and the electrode which loses electrons to the electrolyte is called the anode. The electrical circuit is completed by the movement of ions in solution. Anions move toward the anode, whereas cations move toward the cathode. The cell compartments can be separated by either a porous glass barrier or by a salt bridge to maintain electrical neutrality in the solutions. Galvanic cells must have a positive reduction potential in order for any reaction to occur.

 

Example 2: Copper(ii) - Aluminum cell

Cu(s) | Cu 2+(aq), SO4 2-(aq) || Al 3+(aq), Cl-(aq) | Al (s)

Cathode: 3[Cu2+ + 2e- -->Cu(s)] Eºr = + 0.34V

Anode: 2[Al(s) -->Al3+ + 3e-] Eºr = + 1.66V

Balanced Net Equation: 3Cu2+ + 2Al(s) --> Al3+ + Cu(s)

Eº = +2.00V

The reaction is spontaneous; the copper ions have a stronger pull on the electrons in aluminum than aluminum atoms themselves have, resulting in a reaction. The voltage detected by the voltmeter confirms there is a spontaneous reaction.

Direction of electron flow: anode (-ve) --> cathode (+ve)

 

Descriptions of reactants and products:

Cu(s) - solid, reddish-metalic in colour, malleable

CuSO4- blue translucent solution, odorless

Al(s)- solid, silvery-metalic solid, malleable

AlCl - clear translucent solution, odorless

Product - Copper sulphate solution pales slightly. A positive voltage was detected.

Identification of product: From the net equation, the products are solid copper and Al3+ ions . The paler colour of the copper sulphate solution suggests that the copper ions are being used up.

Initial
Cell layout

Applications – Batteries

A battery is a portable, self-contained electrochemical power source that consists of one or more voltaic cells. For example, the common 1.5-V batteries used to power flashlights and many consumer electronic devices are single voltaic cells. Greater voltages can be achieved by using multiple voltaic cells in a single battery, as in the case in 12-V automotive batteries. Different applications require batteries with different properties. Some common batteries include:

- Lead-Acid Battery (12V): Consists of six voltaic cells in series, each producing 2V. The cathode of each cell consists of lead dioxide packed on a metal grid. The anode of each cell is composed of lead. Both electrodes are immersed in sulphuric acid. One advantage of a lead-acid battery is that it can be recharged. In an automobile the energy necessary for recharging the battery is provided by a generator driven by the engine.

- Alkaline Battery (1.55V): The most common primary battery. The anode of this battery consists of powdered zinc metal immobilized in a gel in contact with a concentrated solution of KOH. The cathode is a mixture of MnO2(s) and graphite, separated from the anode by a porous fabric. The battery is sealed in a steel can to reduce the risk of leakage of the concentrated KOH. The alkaline battery provides far superior performance over the older dry cells that were based on MnO2 and Zn as the electrochemically active substances.

- Nickel-Cadmium, Nickel-Metal Hydride, and Lithium-Ion Batteries The tremendous growth in high-power-demand portable electronic devices, such as cellular phones, notebook computers, and video recorders, has increased the demand for lightweight, readily recharged batteries. One of the most common rechargeable batteries is the Ni-Cad battery. There are drawbacks to Ni-Cad batteries. Cadmium is a toxic heavy metal. Its use increases the weight of batteries and provides an environmental hazard. Some of these problems have been alleviated by the development of NiMH batteries. The newest rechargeable battery to receive large use in consumer electronic devices is the Li-ion battery. It is because Li is a light element, Li-ion batteries can achieve a greater energy density than nickel-based batteries.