Electrochemical cell

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A demonstration electrochemical cell setup resembling the Daniell cell. The two half-cells are linked by a salt bridge carrying ions between them. Electrons flow in the external circuit. ElectrochemCell.png
A demonstration electrochemical cell setup resembling the Daniell cell. The two half-cells are linked by a salt bridge carrying ions between them. Electrons flow in the external circuit.

An electrochemical cell is a device that generates electrical energy from chemical reactions. Electrical energy can also be applied to these cells to cause chemical reactions to occur. [1] Electrochemical cells that generate an electric current are called voltaic or galvanic cells and those that generate chemical reactions, via electrolysis for example, are called electrolytic cells. [2]

Contents

Both galvanic and electrolytic cells can be thought of as having two half-cells: consisting of separate oxidation and reduction reactions.

When one or more electrochemical cells are connected in parallel or series they make a battery. Primary cells are single use batteries.

Types of electrochemical cells

Galvanic cell

A galvanic cell (voltaic cell), named after Luigi Galvani (Alessandro Volta), is an electrochemical cell that generates electrical energy from spontaneous redox reactions. [3]

Galvanic cell with no cation flow Galvanic cell with no cation flow.svg
Galvanic cell with no cation flow

A wire connects two different metals (e.g. zinc and copper). Each metal is in a separate solution; often the aqueous sulphate or nitrate forms of the metal, however more generally metal salts and water which conduct current. [4] A salt bridge or porous membrane connects the two solutions, keeping electric neutrality and the avoidance of charge accumulation. The metal's differences in oxidation/reduction potential drive the reaction until equilibrium. [1]

Key features:

Half cells

Galvanic cells consists of two half-cells. Each half-cell consists of an electrode and an electrolyte (both half-cells may use the same or different electrolytes).

The chemical reactions in the cell involve the electrolyte, electrodes, and/or an external substance (fuel cells may use hydrogen gas as a reactant). In a full electrochemical cell, species from one half-cell lose electrons (oxidation) to their electrode while species from the other half-cell gain electrons (reduction) from their electrode.

A salt bridge (e.g., filter paper soaked in KNO3, NaCl, or some other electrolyte) is used to ionically connect two half-cells with different electrolytes, but it prevents the solutions from mixing and unwanted side reactions. An alternative to a salt bridge is to allow direct contact (and mixing) between the two half-cells, for example in simple electrolysis of water.

As electrons flow from one half-cell to the other through an external circuit, a difference in charge is established. If no ionic contact were provided, this charge difference would quickly prevent the further flow of electrons. A salt bridge allows the flow of negative or positive ions to maintain a steady-state charge distribution between the oxidation and reduction vessels, while keeping the contents otherwise separate. Other devices for achieving separation of solutions are porous pots and gelled solutions. A porous pot is used in the Bunsen cell.

Equilibrium reaction

Each half-cell has a characteristic voltage (depending on the metal and its characteristic reduction potential). Each reaction is undergoing an equilibrium reaction between different oxidation states of the ions: when equilibrium is reached, the cell cannot provide further voltage. In the half-cell performing oxidation, the closer the equilibrium lies to the ion/atom with the more positive oxidation state the more potential this reaction will provide. [1] Likewise, in the reduction reaction, the closer the equilibrium lies to the ion/atom with the more negative oxidation state the higher the potential.

Cell potential

The cell potential can be predicted through the use of electrode potentials (the voltages of each half-cell). These half-cell potentials are defined relative to the assignment of 0 volts to the standard hydrogen electrode (SHE). (See table of standard electrode potentials). The difference in voltage between electrode potentials gives a prediction for the potential measured. When calculating the difference in voltage, one must first rewrite the half-cell reaction equations to obtain a balanced oxidation-reduction equation.

  1. Reverse the reduction reaction with the smallest potential (to create an oxidation reaction/overall positive cell potential)
  2. Half-reactions must be multiplied by integers to achieve electron balance.

Cell potentials have a possible range of roughly zero to 6 volts. Cells using water-based electrolytes are usually limited to cell potentials less than about 2.5 volts due to high reactivity of the powerful oxidizing and reducing agents with water which is needed to produce a higher voltage. Higher cell potentials are possible with cells using other solvents instead of water. For instance, lithium cells with a voltage of 3 volts are commonly available.

The cell potential depends on the concentration of the reactants, as well as their type. As the cell is discharged, the concentration of the reactants decreases and the cell potential also decreases.

Electrolytic cell

An electrolytic cell is an electrochemical cell in which applied electrical energy drives a non-spontaneous redox reaction. [5]

A modern electrolytic cell consisting of two half reactions, two electrodes, a salt bridge, voltmeter, and a battery. Electrolytic Cell Diagram.jpg
A modern electrolytic cell consisting of two half reactions, two electrodes, a salt bridge, voltmeter, and a battery.

They are often used to decompose chemical compounds, in a process called electrolysis. (The Greek word "lysis" (λύσις) means "loosing" or "setting free".)

Important examples of electrolysis are the decomposition of water into hydrogen and oxygen, and of bauxite into aluminium and other chemicals. Electroplating (e.g. of Copper, Silver, Nickel or Chromium) is done using an electrolytic cell. Electrolysis is a technique that uses a direct electric current (DC).

The components of an electrolytic cell are:

When driven by an external voltage (potential difference) applied to the electrodes, the ions in the electrolyte are attracted to the electrode with the opposite potential, where charge-transferring (also called faradaic or redox) reactions can take place. Only with a sufficient external voltage can an electrolytic cell decompose a normally stable, or inert chemical compound in the solution. Thus the electrical energy provided produces a chemical reaction which would not occur spontaneously otherwise.

Key features:

Primary cell

A variety of standard sizes of primary cells. From left: 4.5V multicell battery, D, C, AA, AAA, AAAA, A23, 9V multicell battery, LR44 (top), CR2032 (bottom). Batteries comparison 4,5 D C AA AAA AAAA A23 9V CR2032 LR44 matchstick-1.jpeg
A variety of standard sizes of primary cells. From left: 4.5V multicell battery, D, C, AA, AAA, AAAA, A23, 9V multicell battery, LR44 (top), CR2032 (bottom).

A primary cell produces current by irreversible chemical reactions (ex. small disposable batteries) and is not rechargeable.

They are used for their portability, low cost, and short lifetime.

Primary cells are made in a range of standard sizes to power small household appliances such as flashlights and portable radios.

As chemical reactions proceed in a primary cell, the battery uses up the chemicals that generate the power; when they are gone, the battery stops producing electricity.

Circuit diagram of a primary cell showing difference in cell potential, and flow of electrons through a resistor. Diagram of a primary cell (battery).jpg
Circuit diagram of a primary cell showing difference in cell potential, and flow of electrons through a resistor.

Primary batteries make up about 90% of the $50 billion battery market, but secondary batteries have been gaining market share. About 15 billion primary batteries are thrown away worldwide every year, [6] virtually all ending up in landfills. Due to the toxic heavy metals and strong acids or alkalis they contain, batteries are hazardous waste. Most municipalities classify them as such and require separate disposal. The energy needed to manufacture a battery is about 50 times greater than the energy it contains. [7] [8] [9] [10] Due to their high pollutant content compared to their small energy content, the primary battery is considered a wasteful, environmentally unfriendly technology. Mainly due to the increasing sales of wireless devices and cordless tools, which cannot be economically powered by primary batteries and come with integral rechargeable batteries, the secondary battery industry has high growth and has slowly been replacing the primary battery in high end products.

Secondary cell

Lead acid car battery (secondary cell) Photo-CarBattery.jpg
Lead acid car battery (secondary cell)
Circuit diagram of a secondary cell showing difference in cell potential, and flow of electrons through a resistor. Secondary Cell Diagram.svg
Circuit diagram of a secondary cell showing difference in cell potential, and flow of electrons through a resistor.

A secondary cell produces current by reversible chemical reactions (ex. lead-acid battery car battery) and is rechargeable.

Lead-acid batteries are used in an automobile to start an engine and to operate the car's electrical accessories when the engine is not running. The alternator, once the car is running, recharges the battery.

It can perform as a galvanic cell and an electrolytic cell. It is a convenient way to store electricity: when current flows one way, the levels of one or more chemicals build up (charging); while it is discharging, they reduce and the resulting electromotive force can do work.

They are used for their high voltage, low costs, reliability, and long lifetime.

Fuel cell

Scheme of a proton-conducting fuel cell Solid oxide fuel cell protonic.svg
Scheme of a proton-conducting fuel cell

A fuel cell is an electrochemical cell that reacts hydrogen fuel with oxygen or another oxidizing agent, to convert chemical energy to electricity.

Fuel cells are different from batteries in requiring a continuous source of fuel and oxygen (usually from air) to sustain the chemical reaction, whereas in a battery the chemical energy comes from chemicals already present in the battery.

Fuel cells can produce electricity continuously for as long as fuel and oxygen are supplied.

They are used for primary and backup power for commercial, industrial and residential buildings and in remote or inaccessible areas. They are also used to power fuel cell vehicles, including forklifts, automobiles, buses, boats, motorcycles and submarines.

Fuel cells are classified by the type of electrolyte they use and by the difference in startup time, which ranges from 1 second for proton-exchange membrane fuel cells (PEM fuel cells, or PEMFC) to 10 minutes for solid oxide fuel cells (SOFC).

There are many types of fuel cells, but they all consist of:

anode
At the anode a catalyst causes the fuel to undergo oxidation reactions that generate protons (positively charged hydrogen ions) and electrons. The protons flow from the anode to the cathode through the electrolyte after the reaction. At the same time, electrons are drawn from the anode to the cathode through an external circuit, producing direct current electricity.
cathode
At the cathode, another catalyst causes hydrogen ions, electrons, and oxygen to react, forming water.
electrolyte
Allows positively charged hydrogen ions (protons) to move between the two sides of the fuel cell.

A related technology are flow batteries, in which the fuel can be regenerated by recharging. Individual fuel cells produce relatively small electrical potentials, about 0.7 volts, so cells are "stacked", or placed in series, to create sufficient voltage to meet an application's requirements. [11] In addition to electricity, fuel cells produce water, heat and, depending on the fuel source, very small amounts of nitrogen dioxide and other emissions. The energy efficiency of a fuel cell is generally between 40 and 60%; however, if waste heat is captured in a cogeneration scheme, efficiencies up to 85% can be obtained.

In 2022, the global fuel cell market was estimated to be $6.3 billion, and is expected to increase by 19.9% by 2030. [12] Many countries are attempting to enter the market by setting renewable energy GW goals. [13]

See also

Related Research Articles

<span class="mw-page-title-main">Anode</span> Electrode through which conventional current flows into a polarized electrical device

An anode is an electrode of a polarized electrical device through which conventional current enters the device. This contrasts with a cathode, an electrode of the device through which conventional current leaves the device. A common mnemonic is ACID, for "anode current into device". The direction of conventional current in a circuit is opposite to the direction of electron flow, so electrons flow from the anode of a galvanic cell, into an outside or external circuit connected to the cell. For example, the end of a household battery marked with a "+" is the cathode.

<span class="mw-page-title-main">Cathode</span> Electrode where reduction takes place

A cathode is the electrode from which a conventional current leaves a polarized electrical device such as a lead-acid battery. This definition can be recalled by using the mnemonic CCD for Cathode Current Departs. A conventional current describes the direction in which positive charges move. Electrons have a negative electrical charge, so the movement of electrons is opposite to that of the conventional current flow. Consequently, the mnemonic cathode current departs also means that electrons flow into the device's cathode from the external circuit. For example, the end of a household battery marked with a + (plus) is the cathode.

<span class="mw-page-title-main">Electrochemistry</span> Branch of chemistry

Electrochemistry is the branch of physical chemistry concerned with the relationship between electrical potential difference and identifiable chemical change. These reactions involve electrons moving via an electronically-conducting phase between electrodes separated by an ionically conducting and electronically insulating electrolyte.

<span class="mw-page-title-main">Electrode</span> Electrical conductor used to make contact with nonmetallic parts of a circuit

An electrode is an electrical conductor used to make contact with a nonmetallic part of a circuit. Electrodes are essential parts of batteries that can consist of a variety of materials (chemicals) depending on the type of battery.

<span class="mw-page-title-main">Electrolysis</span> Technique in chemistry and manufacturing

In chemistry and manufacturing, electrolysis is a technique that uses direct electric current (DC) to drive an otherwise non-spontaneous chemical reaction. Electrolysis is commercially important as a stage in the separation of elements from naturally occurring sources such as ores using an electrolytic cell. The voltage that is needed for electrolysis to occur is called the decomposition potential. The word "lysis" means to separate or break, so in terms, electrolysis would mean "breakdown via electricity."

In electrochemistry, electrode potential is the voltage of a galvanic cell built from a standard reference electrode and another electrode to be characterized. By convention, the reference electrode is the standard hydrogen electrode (SHE). It is defined to have a potential of zero volts. It may also be defined as the potential difference between the charged metallic rods and salt solution.

<span class="mw-page-title-main">Redox</span> Chemical reaction in which oxidation states of atoms are changed

Redox is a type of chemical reaction in which the oxidation states of the reactants change. Oxidation is the loss of electrons or an increase in the oxidation state, while reduction is the gain of electrons or a decrease in the oxidation state. The oxidation and reduction processes occur simultaneously in the chemical reaction.

<span class="mw-page-title-main">Galvanic cell</span> Electrochemical device

A galvanic cell or voltaic cell, named after the scientists Luigi Galvani and Alessandro Volta, respectively, is an electrochemical cell in which an electric current is generated from spontaneous oxidation–reduction reactions. A common apparatus generally consists of two different metals, each immersed in separate beakers containing their respective metal ions in solution that are connected by a salt bridge or separated by a porous membrane.

<span class="mw-page-title-main">Electrolytic cell</span> Cell that uses electrical energy to drive a non-spontaneous redox reaction

An electrolytic cell is an electrochemical cell that utilizes an external source of electrical energy to force a chemical reaction that would otherwise not occur. The external energy source is a voltage applied between the cell's two electrodes; an anode and a cathode, which are immersed in an electrolyte solution. This is in contrast to a galvanic cell, which itself is a source of electrical energy and the foundation of a battery. The net reaction taking place in a galvanic cell is a spontaneous reaction, i.e., the Gibbs free energy remains -ve, while the net reaction taking place in an electrolytic cell is the reverse of this spontaneous reaction, i.e., the Gibbs free energy is +ve.

<span class="mw-page-title-main">Standard electrode potential</span> Electromotive force of a half reaction cell versus standard hydrogen electrode

In electrochemistry, standard electrode potential, or , is a measure of the reducing power of any element or compound. The IUPAC "Gold Book" defines it as; "the value of the standard emf of a cell in which molecular hydrogen under standard pressure is oxidized to solvated protons at the left-hand electrode".

A regenerative fuel cell or reverse fuel cell (RFC) is a fuel cell run in reverse mode, which consumes electricity and chemical B to produce chemical A. By definition, the process of any fuel cell could be reversed. However, a given device is usually optimized for operating in one mode and may not be built in such a way that it can be operated backwards. Standard fuel cells operated backwards generally do not make very efficient systems unless they are purpose-built to do so as with high-pressure electrolysers, regenerative fuel cells, solid-oxide electrolyser cells and unitized regenerative fuel cells.

A primary battery or primary cell is a battery that is designed to be used once and discarded, and it is not rechargeable unlike a secondary cell. In general, the electrochemical reaction occurring in the cell is not reversible, rendering the cell unrechargeable. As a primary cell is used, chemical reactions in the battery use up the chemicals that generate the power; when they are gone, the battery stops producing electricity. In contrast, in a secondary cell, the reaction can be reversed by running a current into the cell with a battery charger to recharge it, regenerating the chemical reactants. Primary cells are made in a range of standard sizes to power small household appliances such as flashlights and portable radios.

<span class="mw-page-title-main">Electrolysis of water</span> Electricity-induced chemical reaction

Electrolysis of water is using electricity to split water into oxygen and hydrogen gas by electrolysis. Hydrogen gas released in this way can be used as hydrogen fuel, but must be kept apart from the oxygen as the mixture would be extremely explosive. Separately pressurised into convenient 'tanks' or 'gas bottles', hydrogen can be used for oxyhydrogen welding and other applications, as the hydrogen / oxygen flame can reach approximately 2,800°C.

In electrochemistry, overpotential is the potential difference (voltage) between a half-reaction's thermodynamically determined reduction potential and the potential at which the redox event is experimentally observed. The term is directly related to a cell's voltage efficiency. In an electrolytic cell the existence of overpotential implies that the cell requires more energy than thermodynamically expected to drive a reaction. In a galvanic cell the existence of overpotential means less energy is recovered than thermodynamics predicts. In each case the extra/missing energy is lost as heat. The quantity of overpotential is specific to each cell design and varies across cells and operational conditions, even for the same reaction. Overpotential is experimentally determined by measuring the potential at which a given current density is achieved.

In battery technology, a concentration cell is a limited form of a galvanic cell that has two equivalent half-cells of the same composition differing only in concentrations. One can calculate the potential developed by such a cell using the Nernst equation. A concentration cell produces a small voltage as it attempts to reach chemical equilibrium, which occurs when the concentration of reactant in both half-cells are equal. Because an order of magnitude concentration difference produces less than 60 millivolts at room temperature, concentration cells are not typically used for energy storage.

In electrochemistry, electrosynthesis is the synthesis of chemical compounds in an electrochemical cell. Compared to ordinary redox reactions, electrosynthesis sometimes offers improved selectivity and yields. Electrosynthesis is actively studied as a science and also has industrial applications. Electrooxidation has potential for wastewater treatment as well.

In electrochemistry, Faraday efficiency describes the efficiency with which charge (electrons) is transferred in a system facilitating an electrochemical reaction. The word "Faraday" in this term has two interrelated aspects: first, the historic unit for charge is the faraday (F), but has since been replaced by the coulomb (C); and secondly, the related Faraday's constant correlates charge with moles of matter and electrons. This phenomenon was originally understood through Michael Faraday's work and expressed in his laws of electrolysis.

<span class="mw-page-title-main">Solid oxide electrolyzer cell</span> Type of fuel cell

A solid oxide electrolyzer cell (SOEC) is a solid oxide fuel cell that runs in regenerative mode to achieve the electrolysis of water by using a solid oxide, or ceramic, electrolyte to produce hydrogen gas and oxygen. The production of pure hydrogen is compelling because it is a clean fuel that can be stored, making it a potential alternative to batteries, methane, and other energy sources. Electrolysis is currently the most promising method of hydrogen production from water due to high efficiency of conversion and relatively low required energy input when compared to thermochemical and photocatalytic methods.

The lithium–air battery (Li–air) is a metal–air electrochemical cell or battery chemistry that uses oxidation of lithium at the anode and reduction of oxygen at the cathode to induce a current flow.

Aluminium-ion batteries are a class of rechargeable battery in which aluminium ions serve as charge carriers. Aluminium can exchange three electrons per ion. This means that insertion of one Al3+ is equivalent to three Li+ ions. Thus, since the ionic radii of Al3+ (0.54 Å) and Li+ (0.76 Å) are similar, significantly higher numbers of electrons and Al3+ ions can be accepted by cathodes with little damage. Al has 50 times (23.5 megawatt-hours m-3) the energy density of Li and is even higher than coal.

References

  1. 1 2 3 Wenzel, Thomas J. (July 30, 2013). "Douglas A. Skoog, Donald M. West, F. James Holler, and Stanley R. Crouch: Fundamentals of analytical chemistry, 9th ed., international ed". Analytical and Bioanalytical Chemistry. 405 (25): 412–432. doi:10.1007/s00216-013-7242-1. ISSN   1618-2642. S2CID   94566587.
  2. ↑ Wendt, Hartmut; Kolb, Dieter M.; Engelmann, Gerald E.; Ziegler, Jörg C. (October 15, 2011), "Electrochemistry, 1. Fundamentals", in Wiley-VCH Verlag GmbH & Co. KGaA (ed.), Ullmann's Encyclopedia of Industrial Chemistry, Weinheim, Germany: Wiley-VCH Verlag GmbH & Co. KGaA, pp. a09_183.pub4, doi:10.1002/14356007.a09_183.pub4, ISBN   978-3-527-30673-2 , retrieved May 5, 2023
  3. ↑ Chemistry, Rice University, 2015. [Online]. Available: https://web.ung.edu/media/Chemistry2/Chemistry-LR.pdf
  4. ↑ Ahmad, Dr. Zaki (2013). Principles of corrosion engineering and corrosion control. Butterworth-Heinemann. ISBN   978-0-08-097134-6. OCLC   857524149.
  5. ↑ Brett, C.M.A. (2018), "Standard Electrode Potentials and Application to Characterization of Corrosion Phenomena", Encyclopedia of Interfacial Chemistry, Elsevier, pp. 511–516, doi:10.1016/b978-0-12-409547-2.13389-x, ISBN   9780128098943 , retrieved April 18, 2023
  6. ↑ Communications, Cactus. "What if we could recycle the energy remaining in discarded batteries? Scientists now know how". techxplore.com. Retrieved April 18, 2023.
  7. ↑ Hill, Marquita K. (2004). Understanding Environmental Pollution: A Primer . Cambridge University Press. p. 274. ISBN   978-0-521-82024-0. Manufacturing a disposable battery takes about 50 times more energy than the battery provides when used.
  8. ↑ Watts, John (2006). Gcse Edexcel Science. Letts and Lonsdale. p. 63. ISBN   978-1-905129-63-8.
  9. ↑ Wastebusters Ltd. (2013). The Green Office Manual: A Guide to Responsible Practice. Routledge. p. 96. ISBN   978-1-134-19798-9.
  10. ↑ Danaher, Kevin; Biggs, Shannon; Mark, Jason (2016). Building the Green Economy: Success Stories from the Grassroots. Routledge. p. 199. ISBN   978-1-317-26292-3.
  11. ↑ Qi, Zhaoxiang; Koenig, Gary M. (July 1, 2017). "Review Article: Flow battery systems with solid electroactive materials". Journal of Vacuum Science & Technology B. 35 (4): 040801. Bibcode:2017JVSTB..35d0801Q. doi: 10.1116/1.4983210 . ISSN   2166-2746.
  12. ↑ "Fuel Cell Market Size, Share & Trends Analysis Report, 2030". www.grandviewresearch.com. Retrieved April 18, 2023.
  13. ↑ "Renewable energy targets". energy.ec.europa.eu. Retrieved April 22, 2023.
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