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Half-reaction
Half-reaction
Half-reaction
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In chemistry, a half reaction (or half-cell reaction) is either the oxidation or reduction reaction component of a redox reaction. A half reaction is obtained by considering the change in oxidation states of individual substances involved in the redox reaction. Often, the concept of half reactions is used to describe what occurs in an electrochemical cell, such as a Galvanic cell battery. Half reactions can be written to describe both the metal undergoing oxidation (known as the anode) and the metal undergoing reduction (known as the cathode).

Half reactions are often used as a method of balancing redox reactions. For oxidation-reduction reactions in acidic conditions, after balancing the atoms and oxidation numbers, one will need to add H+ ions to balance the hydrogen ions in the half reaction. For oxidation-reduction reactions in basic conditions, after balancing the atoms and oxidation numbers, first treat it as an acidic solution and then add OH− ions to balance the H+ ions in the half reactions (which would give H2O).

Example: Zn and Cu galvanic cell

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Galvanic cell

Consider the Galvanic cell shown in the adjacent image: it is constructed with a piece of zinc (Zn) submerged in a solution of zinc sulfate (ZnSO4) and a piece of copper (Cu) submerged in a solution of copper(II) sulfate (CuSO4). The overall reaction is:

Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)

At the Zn anode, oxidation takes place (the metal loses electrons). This is represented in the following oxidation half reaction (note that the electrons are on the products side):

Zn(s) → Zn2+ + 2 e−

At the Cu cathode, reduction takes place (electrons are accepted). This is represented in the following reduction half reaction (note that the electrons are on the reactants side):

Cu2+ + 2 e− → Cu(s)

Example: oxidation of magnesium

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Experiment showing synthesis of a basic oxide. Magnesium ribbon is ignited by burner. Magnesium burns emitting intense light and forming magnesium oxide (MgO).
Photograph of a burning magnesium ribbon with very short exposure to obtain oxidation detail.

Consider the example burning of magnesium ribbon (Mg). When magnesium burns, it combines with oxygen (O2) from the air to form magnesium oxide (MgO) according to the following equation:

2 Mg(s) + O2(g) → 2 MgO(s)

Magnesium oxide is an ionic compound containing Mg2+ and O2− ions whereas Mg(s) and O2(g) are elements with no charges. The Mg(s) with zero charge gains a +2 charge going from the reactant side to product side, and the O2(g) with zero charge gains a −2 charge. This is because when Mg(s) becomes Mg2+, it loses 2 electrons. Since there are 2 Mg on left side, a total of 4 electrons are lost according to the following oxidation half reaction:

2 Mg(s) → 2 Mg2+ + 4 e−

On the other hand, O2 was reduced: its oxidation state goes from 0 to −2. Thus, a reduction half reaction can be written for the O2 as it gains 4 electrons:

O2(g) + 4 e− → 2 O2−

The overall reaction is the sum of both half reactions:

2 Mg(s) + O2(g) + 4 e− → 2 Mg2+ + 2 O2− + 4 e−

When chemical reaction, especially, redox reaction takes place, we do not see the electrons as they appear and disappear during the course of the reaction. What we see is the reactants (starting material) and end products. Due to this, electrons appearing on both sides of the equation are canceled. After canceling, the equation is re-written as

2 Mg(s) + O2(g) → 2 Mg2+ + 2 O2−

Two ions, positive (Mg2+) and negative (O2−) exist on product side and they combine immediately to form a compound magnesium oxide (MgO) due to their opposite charges (electrostatic attraction). In any given oxidation-reduction reaction, there are two half reactions—oxidation half reaction and reduction half reaction. The sum of these two half reactions is the oxidation–reduction reaction.

Half-reaction balancing method

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Consider the reaction below:

Cl2 + 2 Fe2+ → 2 Cl− + 2 Fe3+

The two elements involved, iron and chlorine, each change oxidation state; iron from +2 to +3, chlorine from 0 to −1. There are then effectively two half reactions occurring. These changes can be represented in formulas by inserting appropriate electrons into each half reaction:

Fe2+ → Fe3+ + e−
Cl2 + 2 e− → 2 Cl−

Given two half reactions it is possible, with knowledge of appropriate electrode potentials, to arrive at the complete (original) reaction the same way. The decomposition of a reaction into half reactions is key to understanding a variety of chemical processes. For example, in the above reaction, it can be shown that this is a redox reaction in which Fe is oxidised, and Cl is reduced. Note the transfer of electrons from Fe to Cl. Decomposition is also a way to simplify the balancing of a chemical equation. A chemist can atom balance and charge balance one piece of an equation at a time.

For example:

  • Fe2+ → Fe3+ + e− becomes 2 Fe2+ → 2 Fe3+ + 2e−
  • is added to Cl2 + 2 e− → 2 Cl−
  • and finally becomes Cl2 + 2 Fe2+ → 2 Cl− + 2 Fe3+

It is also possible and sometimes necessary to consider a half reaction in either basic or acidic conditions, as there may be an acidic or basic electrolyte in the redox reaction. Due to this electrolyte it may be more difficult to satisfy the balance of both the atoms and charges. This is done by adding H2O, OH−, e−, and/or H+ to either side of the reaction until both atoms and charges are balanced.

Consider the half reaction below:

PbO2 → PbO

OH−, H2O, and e− can be used to balance the charges and atoms in basic conditions, as long as it is assumed that the reaction is in water.

2 e− + H2O + PbO2 → PbO + 2 OH−

Again consider the half reaction below:

PbO2 → PbO

H+, H2O, and e− can be used to balance the charges and atoms in acidic conditions, as long as it is assumed that the reaction is in water.

2 e− + 2 H+ + PbO2 → PbO + H2O

Notice that both sides are both charge balanced and atom balanced.

Often there will be both H+ and OH− present in acidic and basic conditions but that the resulting reaction of the two ions will yield water.

H+ + OH− → H2O

See also

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References

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Revisions and contributorsEdit on WikipediaRead on Wikipedia

Half-reaction

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from Grokipedia
A half-reaction is a chemical equation that describes either the oxidation or reduction component of a redox (oxidation-reduction) reaction, explicitly showing the transfer of electrons while isolating one half of the overall process. In a complete redox reaction, two half-reactions—one for oxidation (loss of electrons) and one for reduction (gain of electrons)—combine to form the full equation, ensuring conservation of mass and charge. This separation is fundamental to understanding electron transfer in chemical systems, as it allows for the independent analysis of each process. Half-reactions are essential for balancing equations, particularly in acidic or basic solutions, using a systematic method that first balances atoms other than and oxygen, then adjusts for oxygen with , with protons (or in basic media), and finally charge with electrons. For instance, the oxidation half-reaction for might be written as Zn → Zn²⁺ + 2e⁻, while the reduction of (II) ions is Cu²⁺ + 2e⁻ → Cu. This approach ensures that the number of electrons lost in oxidation equals those gained in reduction when combining the halves. In , half-reactions occur at electrodes in electrochemical cells, such as galvanic cells where spontaneous processes generate electricity: oxidation at the and reduction at the , with flowing through an external circuit. Standard reduction potentials, measured relative to the hydrogen electrode (assigned 0 ), quantify the tendency of a to gain and predict reaction spontaneity; for example, the standard potential for Zn²⁺/Zn is -0.76 , indicating a strong . These potentials are critical in applications ranging from batteries to biological electron transport chains.

Fundamentals

Definition

A half-reaction represents either the oxidation or reduction component of a complete redox reaction, obtained by isolating the changes in oxidation states and balancing the equation with electrons to account for charge differences. This approach simplifies the analysis of redox processes by decoupling the electron transfer mechanisms that occur simultaneously in the full reaction, allowing chemists to study each part independently, particularly in electrochemical contexts. In an oxidation half-reaction, electrons are lost as the of the species increases, with the serving as the . Conversely, a reduction half-reaction involves the gain of electrons, decreasing the , where the accepts the electrons. These distinctions highlight the complementary nature of the two half-reactions, which together form a balanced pair without net electron accumulation. Half-reactions are conventionally notated using standard formats, with reactants on the left and products on the right separated by an arrow indicating direction. For oxidation, electrons appear as products (e.g., reactant → product + e⁻), while for reduction, they are reactants (e.g., reactant + e⁻ → product), ensuring charge balance in each isolated equation. The conceptual foundations of half-reactions trace to 18th-century advancements in understanding processes, pioneered by , who reframed and as oxidation reactions involving oxygen rather than the release of phlogiston. 's work in the and established the dualistic view of oxidation and reduction, laying the groundwork for later electrochemical interpretations that formalized these as separable half-reactions.

Role in Redox Processes

Half-reactions serve as the fundamental building blocks of processes, isolating the oxidation and reduction components of a complete reaction. The oxidation half-reaction depicts the loss of electrons by a species, while the reduction half-reaction shows the gain of electrons by another. To form the overall reaction, these half-reactions are combined by multiplying them to equalize the number of electrons transferred, then adding them together, with the electrons canceling out. This approach simplifies the of complex reactions by allowing chemists to examine each process independently before integration. In electrochemical systems, half-reactions are essential for understanding and predicting the behavior of voltaic cells, where the oxidation half-reaction occurs at the and the reduction half-reaction at the . The standard cell potential, Ecell∘E^\circ_\text{cell}, is calculated as Ecell∘=Ecathode∘−Eanode∘E^\circ_\text{cell} = E^\circ_\text{cathode} - E^\circ_\text{anode}, using tabulated standard reduction potentials (E∘E^\circ) for each half-reaction. A positive Ecell∘E^\circ_\text{cell} indicates a spontaneous reaction, enabling the design and optimization of devices like batteries. This method allows initial assessment of reaction feasibility without deriving the full coupled equation. Standard half-reaction potentials are referenced to the standard hydrogen electrode (SHE), defined by the half-reaction 2H+(aq,1 M)+2e−⇌H2(g,1 bar)2\text{H}^+ (\text{aq}, 1~\text{M}) + 2\text{e}^- \rightleftharpoons \text{H}_2 (\text{g}, 1~\text{bar}) with E∘=0 VE^\circ = 0~\text{V} at 25°C. These values are tabulated for common species, such as Cu2++2e−⇌Cu\text{Cu}^{2+} + 2\text{e}^- \rightleftharpoons \text{Cu} at +0.337 V, facilitating comparisons and predictions across diverse systems. Half-reactions are particularly valuable in analyzing processes like corrosion, where the anodic oxidation of a metal (e.g., Fe→Fe2++2e−\text{Fe} \rightarrow \text{Fe}^{2+} + 2\text{e}^-) couples with a cathodic reduction (e.g., oxygen or hydrogen evolution), without immediately requiring the complete equation. Similarly, in batteries, they underpin the evaluation of energy storage and discharge mechanisms, aiding in the development of efficient electrochemical technologies.

Examples

Zinc-Copper Galvanic Cell

The Daniell cell, a prototypical galvanic cell, features a zinc anode immersed in a zinc sulfate (ZnSO₄) solution and a copper cathode immersed in a copper(II) sulfate (CuSO₄) solution, with the two half-cells separated by a porous barrier or salt bridge to prevent direct mixing while allowing ion migration. At the , oxidation occurs according to the half-reaction:
\ceZn(s)−>Zn2+(aq)+2e−\ce{Zn(s) -> Zn^{2+}(aq) + 2e^-}
This process releases electrons that flow through the external circuit toward the . At the , reduction takes place via the half-reaction:
\ceCu2+(aq)+2e−−>Cu(s)\ce{Cu^{2+}(aq) + 2e^- -> Cu(s)}
Here, copper(II) ions in solution accept the electrons to deposit as copper metal on the . The overall cell reaction combines these half-reactions:
\ceZn(s)+Cu2+(aq)−>Zn2+(aq)+Cu(s)\ce{Zn(s) + Cu^{2+}(aq) -> Zn^{2+}(aq) + Cu(s)}
Under standard conditions, the cell potential is calculated as E\ce∘cell=E\ce∘Cu2+/Cu−E\ce∘Zn2+/Zn=0.34 V−(−0.76 V)=1.10 VE^\circ_\ce{cell} = E^\circ_\ce{Cu^{2+}/Cu} - E^\circ_\ce{Zn^{2+}/Zn} = 0.34 \, \text{V} - (-0.76 \, \text{V}) = 1.10 \, \text{V}, indicating a spontaneous process that generates electrical energy.

Magnesium Oxidation

Magnesium demonstrates significant reactivity in aqueous environments, undergoing oxidation when reacting with or acids to liberate gas. In acidic conditions, such as with , magnesium metal displaces ions, resulting in a complete process where magnesium is oxidized and protons are reduced. This reaction is notably vigorous, producing due to the rapid evolution of gas bubbles, and is exothermic, releasing heat that can cause the solution temperature to rise appreciably. The oxidation half-reaction for magnesium is represented as: \ceMg(s)−>Mg2+(aq)+2e−\ce{Mg(s) -> Mg^2+(aq) + 2e^-} This process involves the loss of two electrons per magnesium atom, converting the neutral metal to the divalent cation in solution. For contextual completeness in acidic media, the corresponding reduction half-reaction is: \ce2H+(aq)+2e−−>H2(g)\ce{2H+(aq) + 2e^- -> H2(g)} However, the focus remains on the oxidation of magnesium, which drives the overall reactivity observed. With alone, the reaction proceeds more slowly, particularly at , forming magnesium hydroxide and gas, but accelerates under heated conditions. Beyond laboratory demonstrations, the oxidative behavior of magnesium finds practical application in corrosion prevention through sacrificial anodes. In this method, magnesium, being more reactive than metals like iron or , is connected to the protected structure, preferentially oxidizing and thereby cathodically protecting the less active metal from in environments such as or . This sacrificial role extends the lifespan of pipelines, ship hulls, and underground infrastructure by continuously supplying electrons to inhibit anodic dissolution of the primary material.

Balancing Procedures

Acidic Conditions

Balancing half-reactions in acidic conditions follows a systematic procedure that accounts for the availability of hydrogen ions (H⁺) in the medium, ensuring conservation of both mass and charge. This method is essential for preparing half-reactions that can be combined to form complete equations, aiding in the prediction of cell potentials in electrochemical cells. The procedure consists of four main steps:
  1. Balance all elements except oxygen and hydrogen: Adjust coefficients to equalize the number of atoms for all species other than O and H on both sides of the half-reaction equation.
  2. Balance oxygen using H₂O: Add water molecules to the side deficient in oxygen atoms to achieve equality.
  3. Balance hydrogen using H⁺: Add hydrogen ions to the side deficient in hydrogen atoms, leveraging the acidic environment.
  4. Balance charge by adding e⁻: Calculate the total charge on each side and add electrons to the more positive side (for reductions, electrons are reactants; for oxidations, products) to equalize charges.
To illustrate, consider the reduction half-reaction for permanganate ion to manganese(II) ion in acidic solution: MnO₄⁻ → Mn²⁺.
  • Step 1: is already balanced (1 on each side).
  • Step 2: Add 4 H₂O to the right to balance the 4 oxygen atoms on the left.
  • Step 3: Add 8 H⁺ to the left to balance the 8 hydrogen atoms from the 4 H₂O.
  • Step 4: The left side now has a charge of +7 (from MnO₄⁻ and 8 H⁺), while the right has +2; add 5 e⁻ to the left to balance the charge.
The balanced equation is: MnO4−+8H++5e−→Mn2++4H2O\text{MnO}_4^- + 8\text{H}^+ + 5\text{e}^- \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O} Common pitfalls in this process include adding electrons before completing the atom balances, which can lead to inconsistencies, and neglecting to verify both atomic and charge equality after all steps. Always confirm the final half-reaction by recounting atoms and charges on both sides to ensure accuracy.

Basic Conditions

To balance half-reactions in basic conditions, first follow the procedure for acidic media by balancing all atoms other than hydrogen and oxygen, then adding H₂O to balance oxygen and H⁺ to balance hydrogen, and finally adding electrons to balance charge. After achieving balance in this manner, add an equal number of OH⁻ ions to both sides of the equation to neutralize the H⁺ ions, converting them to H₂O on the side originally containing H⁺ while leaving excess OH⁻ on the opposite side; cancel any common H₂O molecules that appear on both sides. This step eliminates free protons, adapting the equation to the absence of acidic species in basic solution. For instance, the reduction half-reaction of permanganate ion to in basic solution, initially balanced in acidic form as MnO₄⁻ + 4H⁺ + 3e⁻ → MnO₂ + 2H₂O, is adjusted by adding 4OH⁻ to both sides, yielding: MnO4−+2H2O+3e−→MnO2+4OH−\mathrm{MnO_4^- + 2H_2O + 3e^- \rightarrow MnO_2 + 4OH^-} Verification involves confirming that all atoms except H and O are balanced, that the total charge is equal on both sides, and that H and O atoms are accounted for through H₂O and OH⁻ without unpaired H⁺. This approach accurately represents processes in alkaline environments, such as alkaline batteries using KOH or biological systems at near-neutral to basic pH where enzymes facilitate . In variations where the half-reaction occurs inherently in basic media without acidic intermediates, balance directly by using OH⁻ for hydrogen adjustments and H₂O for oxygen, adding electrons for charge neutrality, which avoids the initial acidic balancing step.

References

  1. https://www2.chem.wisc.edu/deptfiles/genchem/netorial/rottosen/tutorial/modules/electrochemistry/02half_reactions/18_21.htm
  2. https://greenchemistry.yale.edu/modrnu-modules/physicochemical-properties-modules/module-2-redox-reactions
  3. https://books.byui.edu/general_college_chemistry_2/redox_electrochemica
  4. https://chem.libretexts.org/Bookshelves/Analytical_Chemistry/Supplemental_Modules_%28Analytical_Chemistry%29/Electrochemistry/Redox_Chemistry/Half-Reactions
  5. https://www.corrosionpedia.com/definition/1728/half-cell-reaction
  6. https://chemed.chem.purdue.edu/genchem/topicreview/bp/ch20/electro.php
  7. https://www.acs.org/education/whatischemistry/landmarks/lavoisier.html
  8. https://chemed.chem.purdue.edu/genchem/history/lavoisier.html
  9. https://pressbooks-dev.oer.hawaii.edu/chemistry/chapter/standard-reduction-potentials/
  10. https://www.usna.edu/NAOE/_files/documents/Courses/EN380/Course_Notes/Corrosion%20Basics.pdf
  11. https://chem.libretexts.org/Bookshelves/General_Chemistry/ChemPRIME_%28Moore_et_al.%29/17%253A_Electrochemical_Cells/17.08%253A_Galvanic_Cells
  12. https://byjus.com/chemistry/daniell-cell/
  13. https://www.vedantu.com/chemistry/daniell-cell
  14. https://chem.libretexts.org/Bookshelves/Analytical_Chemistry/Supplemental_Modules_%28Analytical_Chemistry%29/Electrochemistry/Redox_Potentials/Standard_Potentials
  15. http://hyperphysics.phy-astr.gsu.edu/hbase/Tables/electpot.html
  16. https://chem.libretexts.org/Bookshelves/Introductory_Chemistry/Introductory_Chemistry_%28CK-12%29/23%253A_Electrochemistry/23.06%253A_Calculating_Standard_Cell_Potentials
  17. https://sites.lib.jmu.edu/chemdemos/2013/06/11/bubbling-magnesium/
  18. https://www.colorado.edu/lab/lecture-demo-manual/general-reactions/s115-stoichiometry-limiting-reagents-mg-hcl
  19. https://www.chm.uri.edu/weuler/chm112/lectures/Mgox-13.html
  20. https://www.colorado.edu/lab/lecture-demo-manual/electrochemistry/x825-orange-juice-clock
  21. https://van.physics.illinois.edu/ask/listing/14978
  22. https://chem.washington.edu/lecture-demos/cathodic-protection
  23. https://user.engineering.uiowa.edu/~swan/courses/53086/Corrosion_Protection.pdf
  24. https://chem.libretexts.org/Bookshelves/Analytical_Chemistry/Supplemental_Modules_%28Analytical_Chemistry%29/Electrochemistry/Redox_Chemistry/Balancing_Redox_reactions
  25. https://www.khanacademy.org/science/ap-chemistry-beta/x2eef969c74e0d802:chemical-reactions/x2eef969c74e0d802:oxidation-reduction-redox-reactions/v/balance-and-redox-reactions1
  26. https://www.ck12.org/flexi/chemistry/Balancing-Redox-Reactions-Half-Reaction-Method/write-a-balanced-half-reaction-for-the-reduction-of-mno4-to-mn2plus-in-an-acidic-solution-mno4-%28aq%29-plus-4-h2o%28l%29-greater-mn2plus%28aq%29-plus-8-oh-%28aq%29-mno4-%28aq%29-plus-4-h2o%28l%29-plus-5-e-greater-mn2plus%28aq%29-plus-8-oh-%28aq%29-mno4-%28aq%29-gre/
  27. https://www.chemteam.info/Redox/Balance-Redox-Acid.html
  28. https://tigerweb.towson.edu/ladon/redoxb.html
  29. https://www.tsc.fl.edu/media/divisions/learning-commons/resources-by-subject/science/chemistry/chm1045-amp-l/BalancingRedoxReactions.pdf
  30. https://lavelle.chem.ucla.edu/wp-content/supporting-files/Chem14B/Balancing_Redox_Reactions_Basic_Conditions.pdf
  31. https://mccord.cm.utexas.edu/chembook/page.php?chnum=7&sect=9
  32. https://wou.edu/chemistry/courses/online-chemistry-textbooks/ch103-allied-health-chemistry/ch103-chapter-6-introduction-to-organic-chemistry-and-biological-molecules/
  33. https://uh.edu/~chembi/chem1332ch18.pdf
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