Name: electrolysis in batteries and metal production
Uploaded: 2020-09-24T15:06:52.000Z
Duration: 3 min
Channel: JoVE

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Electrolysis

Spontaneous redox reactions in voltaic cells convert chemical energy into electrical energy. Alternatively, external electrical energy drives nonspontaneous redox reactions in a process called electrolysis, taking place in an electrolytic cell.
An electrolytic cell has an anode for oxidation and a cathode for reduction. However, unlike the voltaic cell, where the negatively charged anode is the source of electrons, the electrolytic cell has an external source of electrons like a battery. 
The battery draws electrons from the anode connected to its positive terminal and drives electrons towards the cathode, which is connected to its negative terminal. Thus electrons flow from the positive to the negative terminal of the external power source.
Electrolysis is used to separate ionic compounds into their constituent elements. In their molten or aqueous state, the constituent ions are free to conduct electricity.
Consider potassium bromide that melts at a temperature of 734 &#176;C. Here, bromide and potassium ions are present at their lowest or highest oxidation state, respectively. Under power, the bromide ions are oxidized to bromine gas at the anode, and potassium ions are reduced to solid potassium at the cathode. 
In molten ionic mixtures like sodium iodide and potassium bromide, standard electrode potentials are used to predict which substance will react. Generally, anions with a lower electrode potential are oxidized, and cations with a more positive electrode potential are reduced.
Interestingly, electrolysis of aqueous potassium bromide produces different products from molten salts due to the possible electrolysis of water.
Thus, two reactions could occur at each electrode: the oxidation of bromide ions or water at the anode and the reduction of potassium ions or water at the cathode.
Given the lower electrode potential, oxidation of bromide ions to bromine is favored, while the more positive electrode potential of water drives the reduction to hydrogen gas. 
Stoichiometry of the half-reactions can be used to calculate the amount of substances produced or consumed during electrolysis.
For example, electrolysis of molten potassium bromide uses one electron to form a potassium atom. Thus, every mole of electrons, or 96,485 coulombs of charge, passed through the cell produces one mole or nearly 39 grams of solid potassium.

electrolysis

In a galvanic cell, the electrical work is done by a redox system on its surroundings as electrons produced by the spontaneous redox reactions are transferred through an external circuit. Alternatively, an external circuit does work on a redox system by imposing a voltage sufficient to drive an otherwise nonspontaneous reaction in a process known as electrolysis. For instance, recharging a battery involves the use of an external power source to drive the spontaneous (discharge) cell reaction in the reverse direction, restoring to some extent the composition of the half-cells and the voltage of the battery. Other examples include the use of electrolysis in the refinement of metallic ores, the manufacture of commodity chemicals, and the electroplating of metallic coatings onto utensils, jewelry, etc.
Predicting the Product of Electrolysis
The electrolysis of molten sodium chloride, NaCl (l), is used for the industrial production of metallic sodium, Na, and chlorine gas, Cl2. Sodium ions (Na+) are reduced to atoms at the cathode, while chloride ions (Cl&#8722;) ions are oxidized to chlorine gas, Cl2, at the anode. The redox reactions are:
<img alt="Electric dipole configuration; diagram with charge \(q\) separation \(d\) and dipole moment direction." src="/files/ftp_upload/11441/11441_Equation_1.png" style="text-align: center; height: 110px;" data-altupdatedat="1761830439000" data-alt="Eq1" />
The negative cell potential indicates a nonspontaneous reaction that must be driven by imposing a positive potential of magnitude greater than &#8722;4.07 V. Electrolysis of molten sodium chloride is carried out at a high temperature of 801 &#176;C as ionic solids melt at high temperatures.
The electrolysis of water produces stoichiometric amounts of oxygen gas at the anode and hydrogen at the cathode. To improve electrical conductivity, the hydrogen ion concentration of the water is increased by adding a strong acid. The associated redox reactions are:
<img alt="Chemical equilibrium; reaction equation CH4 + 2O2 = CO2 + 2H2O; combustion process diagram." src="/files/ftp_upload/11441/11441_Equation_2.png" style="text-align: center; height: 70px;" data-altupdatedat="1761830439000" data-alt="Eq2" />
The electrolysis of solutions of ionic compounds such as aqueous sodium chloride may involve the electrolysis of either water species (H2O, H+, OH-) or solute species (the cations and anions of the compound) at the anode and cathode.
The electrolysis of aqueous sodium chloride could involve either of these two anode reactions:
<img alt="Chemical equilibrium equations, &#931;MA=0, formula diagram, statics education, reaction dynamics." src="/files/ftp_upload/11441/11441_Equation_3.png" style="text-align: center; height: 70px;" data-altupdatedat="1761830439000" data-alt="Eq3" />
Thermodynamically, water would be more readily oxidized; however, in practice, chlorine gas is produced. Oxidation of water would require a much larger voltage to initiate. To overcome this overvoltage, electrodes are chosen and the potential of the cell is carefully monitored to ensure that the oxidation of chloride ions at the anode.
Similarly, the possible reduction reactions at the cathode are:
<img alt="DNA replication process diagram with leading and lagging strands, highlighting Okazaki fragments." src="/files/ftp_upload/11441/11441_Equation_4.png" style="text-align: center; height: 70px;" data-altupdatedat="1761830439000" data-alt="Eq4" />
Comparing these standard half-reaction potentials suggests the reduction of water is thermodynamically favored. The net cell reaction in this case is then:
<img alt="Photosynthesis, light-dependent reactions diagram; chlorophyll, ATP synthesis, electron transport chain." src="/files/ftp_upload/11441/11441_Equation_5.png" style="text-align: center; height: 50px;" data-altupdatedat="1761830439000" data-alt="Eq5" />
Quantitative Aspects of Electrolysis
Electrical current is the rate of flow of electrons and is measured in ampere, one coulomb per second (A = 1 C/s). The charge transferred, Q, by passage of a constant current, I, over a specified time interval, t, is given by
<img alt="Equilibrium equations diagram, &#931;Fx=0, &#931;Fy=0, static balance analysis, physics study." src="/files/ftp_upload/11441/11441_Equation_6.png" style="text-align: center; height: 50px;" data-altupdatedat="1761830439000" data-alt="Eq6" />
When electrons are transferred during a redox process, the stoichiometry of the reaction may be used to derive the total amount of (electronic) charge involved. For example, the generic reduction process,
<img alt="Equilibrium diagram with &#931;Fx=0, &#931;Fy=0, depicting static equilibrium conditions in mechanics." src="/files/ftp_upload/11441/11441_Equation_7.png" style="text-align: center; height: 50px;" data-altupdatedat="1761830439000" data-alt="Eq7" />
involves the transfer of n mole of electrons. The charge transferred is, therefore,
<img alt="Static equilibrium diagram, hydrostatic system, &#931;Fy=0, forces labeled, physics concept." src="/files/ftp_upload/11441/11441_Equation_8.png" style="text-align: center; height: 50px;" data-altupdatedat="1761830439000" data-alt="Eq8" />
where F is Faraday&#8217;s constant, the charge in coulombs for one mole of electrons. For an electrochemical cell, the current flow is measured and can be used in stoichiometric calculations related to the cell reaction.
This text is adapted from <a href="https://openstax.org/books/chemistry-2e/pages/17-7-electrolysis">OpenStax, Chemistry 2e, Section 17.7: Electrolysis</a>.

In a galvanic cell, the electrical work is done by a redox system on its surroundings as electrons produced by the spontaneous redox reactions are ...

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KEY TERMS AND CONCEPTS

electrolysis in batteries and metal production

Electrolysis uses electrical energy to drive a nonspontaneous redox reaction. In a galvanic cell, a spontaneous redox reaction does electrical work on the surroundings by sending electrons through an external circuit. In electrolysis, an external circuit does work on the redox system by applying a voltage large enough to force the reaction in the opposite direction.

This process is used in several real-world settings. Recharging a battery is one example, because an outside power source drives the discharge reaction backward and helps restore the composition of the half-cells and the battery voltage. Electrolysis is also used to refine metallic ores, make commodity chemicals, and plate thin metallic coatings onto utensils and jewelry.

The product of electrolysis depends on the substance being electrolyzed. Molten sodium chloride, NaCl(l), is used to make metallic sodium and chlorine gas. At the cathode, sodium ions are reduced to sodium atoms. At the anode, chloride ions are oxidized to chlorine gas. This reaction has a negative cell potential, so it is nonspontaneous and must be driven by a positive potential greater than 4.07 V. Because sodium chloride is an ionic solid, this electrolysis is carried out at 801 °C, where it melts.

Electrolysis of water produces oxygen gas at the anode and hydrogen gas at the cathode. The water is made more conductive by adding a strong acid, which increases the hydrogen ion concentration. Solutions of ionic compounds, such as aqueous sodium chloride, can be more complex because either water species or the dissolved ions may react at the anode and cathode.

For aqueous sodium chloride, the anode could oxidize either water or chloride ions. Thermodynamically, water is more easily oxidized, but in practice chlorine gas is produced. Oxidation of water needs a much larger starting voltage. To overcome this overvoltage, the electrodes are chosen carefully and the cell potential is monitored so chloride ions are oxidized instead. At the cathode, the possible reduction reactions also include water species and the sodium ions. Comparing the standard reduction potentials shows that water reduction is favored, and the overall cell reaction follows from those half-reactions.

Electrolysis can also be treated quantitatively. Electrical current is the rate of electron flow and is measured in amperes, where 1 A equals 1 C/s. The charge transferred, Q, equals current times time for a constant current. When electrons are part of a redox reaction, the stoichiometry gives the amount of charge transferred in moles of electrons. Faraday’s constant gives the charge for one mole of electrons. Measured current can then be used for stoichiometric calculations in an electrochemical cell.

Electrolysis in Batteries and Metal Production

electrolysis-in-batteries-and-metal-production

balancing redox equations

Redox Reactions and Half-Reaction Balancing

Redox reactions involve electron transfer and changes in oxidation number. In electrochemistry, these reactions connect chemical changes with electrical energy. They also power technologies such as industrial metal processing and rechargeable batteries used in electric vehicles. Oxidation number is a way to track how an atom's electronic environment changes compared with the pure element, where the oxidation number is zero. For a simple ionic compound, the oxidation number matches the ion...

electromotive force

Cell Voltage and Electron Flow

Electric current is the flow of electric charge through a conducting medium or a solution. In wires, that charge is carried by electrons. In some solutions, ions can also move and carry charge. When electrons move through a wire, they produce an electric current. This idea connects to redox reactions, where electrons are lost and gained. In the spontaneous reaction between zinc and copper ions, zinc is oxidized to zinc ions. Copper ions are reduced to solid copper. Even so, this direct...

voltaic/galvanic cells

Galvanic Cells and Battery Reactions

Galvanic cells use spontaneous redox reactions to produce electric current. A common example is a disposable AA battery, which powers devices by turning chemical energy into electrical energy. Another example is a copper wire placed in silver nitrate solution, where the solution slowly turns bright blue and a grey solid forms on the wire. In the copper and silver reaction, copper is oxidized and silver ions are reduced. Copper forms cuprous ions, which give the solution its blue color. At the...

standard electrode potentials

Using Electrode Potentials to Predict Reactions

Electrode potentials help predict whether a redox reaction will happen on its own. A redox reaction involves electron transfer. In electrochemistry, the measured value is called cell potential or cell voltage. It shows the energy linked to charge transfer and is measured in volts, where 1 volt equals 1 joule per coulomb. A single half-cell cannot be measured by itself because electron transfer needs both a donor and a receiver. For that reason, scientists measure the difference between two...

cell potential and free energy

Redox Cell Potential, ΔG°, and K

Redox cell potential connects chemical energy, free energy change, and the equilibrium constant in an electrochemical cell. In these reactions, chemical energy is converted into electrical energy. That link helps predict whether a reaction is spontaneous and how far it may proceed. Thermodynamics is the study of how heat and other forms of energy are related. In an electrochemical cell, cell potential can also reflect oxidizing strength or reducing strength. A similar idea appears in acid-base...

the nernst equation

Cell Potential Under Changing Conditions

Cell potential changes when a redox reaction does not run at standard conditions. Standard cell potentials connect with standard free energy change, ΔG°, and the equilibrium constant K, but those relationships only hold at 1 M concentration, 298 K, and 1 atmosphere. A zinc(II) and tin(II) redox cell with E°cell = +0.291 V and ΔG° = −56.2 kJ is spontaneous under those standard conditions. As a cell discharges, the reactant concentration changes and the cell potential steadily drops. Many redox...

concentration cells

Cell Potential from Concentration Differences

Concentration cells are voltaic cells that use two almost identical half-cells. The half-cells use the same electrode and the same half-reaction. They differ only in the concentration of one redox species, so the cell potential depends on that concentration difference. In a concentration cell, the Nernst equation can be used to find the cell potential. For the zinc example shown, the calculated cell potential is positive. A positive cell potential means the overall reaction is spontaneous. <br...

batteries and fuel cells

Primary, Rechargeable Batteries and Fuel Cells

Batteries and fuel cells are galvanic cells that supply electrical power for many applications. They range from tiny button batteries in wristwatches to large systems that can provide backup energy to municipal power grids. Some are primary cells, which are single-use and cannot be recharged. Others are secondary cells, which use reversible reactions and can be charged again by an external power source. A common primary battery is the dry cell. It uses zinc as both the container and the anode,...

corrosion

Rusting and Metal Protection

Corrosion is the natural breakdown of metals through electrochemical reactions. It can appear as rust on iron, tarnish on silver, or a blue-green patina on copper. In each case, the metal is oxidized. Some corrosion products can protect the surface, while others keep damaging the metal. Rusting is the harmful type of corrosion seen in iron. It happens when iron is exposed to oxygen and water. A galvanic cell forms on the iron surface, and this drives oxidation at the anodic region and...

Electrolysis uses electrical energy to drive a nonspontaneous redox reaction. In a galvanic cell, a spontaneous redox reaction does electrical work on the surroundings by sending electrons through an external circuit. In electrolysis, an external circuit does work on the redox system by applying a voltage large enough to force the reaction in the opposite direction. This process is used in several real-world settings. Recharging a battery is one example, because an outside power source drives...

key terms and concepts

Learn how electrolysis drives nonspontaneous redox reactions, including battery recharging, sodium chloride electrolysis, water splitting, and Faraday-based calculations.

Electrolysis in Batteries and Metal Production | JoVE K12

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Electrolysis in Batteries and Metal Production

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