Name: redox reactions and half-reaction balancing
Uploaded: 2020-09-24T15:06:52.000Z
Duration: 2 min 58 s
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Matter and its Interactions

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Balancing Redox Equations

How does a battery provide electricity to operate portable devices? Why does a cut avocado turn brown, or metal rust underwater? 
These processes are driven by specific types of chemical reactions involving electron transfer from one atom to another. The atom losing electrons is oxidized, while the atom gaining electrons is reduced. These reactions are called oxidation-reduction or redox reactions and are characterized by changes in the oxidation state for one or more reactants. 
The interaction between magnesium and concentrated hydrochloric acid is an example of a redox reaction. Here, magnesium is oxidized to magnesium 2+ ions, and the protons are reduced to hydrogen gas.
Simple chemical equations can be easily balanced. However, as most redox equations are complex, the number of electrons lost and gained must be justified.
When balancing redox equations, it is essential to observe the law of conservation of mass. Each element&#8217;s amount and any gain or loss of electrons must be balanced on both sides of the reaction; thus, if one reactant is reduced, another must be oxidized.
The method of half-reactions is used to balance redox equations successfully.
Consider the reaction of permanganate with oxalate. For aqueous acidic solutions, assign the oxidation states and divide the reaction into its component half-reactions.
Then, balance each half-reaction omitting hydrogen and oxygen. Here, manganese is already balanced, but the oxalate needs a coefficient of two; secondly, balance oxygen atoms with the addition of water molecules; and thirdly, balance hydrogen atoms by adding protons where required.
Balance the charges by adding electrons. Since manganese is reduced, five electrons are added on the reactant side. Conversely, oxalate is oxidized; thus, two electrons are added on the product side. Multiply the half-reactions by an integer so the number of electrons added is equal.
Lastly, add and simplify the balanced half-reactions by canceling out species on both sides to obtain a balanced redox reaction.
For basic aqueous solutions, the procedure is similar but includes one additional step. Here, an equal number of hydroxyl ions is added to both sides of the reaction to neutralize the protons before summating the balanced half-reactions.

balancing redox equations

Electrochemistry is the science involved in the interconversion of electrical and chemical reactions. Such reactions are called reduction-oxidation, or redox reactions. These important reactions are defined by changes in oxidation states for one or more reactant elements and include a subset of reactions involving the transfer of electrons between reactant species. Electrochemistry as a field has evolved to yield sufficient insights on the fundamental principles of redox chemistry and multiple technologies ranging from industrial-scale metallurgical processes to robust, rechargeable batteries for electric vehicles. As reactions involving the transfer of electrons are essential to the study of electrochemistry, a brief review of redox chemistry includes the following.
Oxidation Numbers
By definition, a redox reaction involves a change in either oxidation number or oxidation state for one or multiple reactants. The oxidation number of an element is an assessment of how the electronic environment of its atoms differs compared to atoms of the pure element. By this definition, an atom in an element carries an oxidation number of zero. For an atom, the oxidation number is equal to the atom's charge in the compound if the compound were ionic. Thus, the sum of oxidation numbers for all atoms in a molecule is equal to the molecule's charge.
Ionic Compounds
Simple ionic compounds are the simplest examples of this formalism since the elements have oxidation numbers equal to their ionic charges. Sodium chloride, NaCl, is composed of Na+ cations and Cl&#8722; anions, with oxidation numbers for sodium and chlorine being +1 and &#8722;1, respectively. Calcium fluoride, CaF2, is composed of Ca2+ cations and F&#8722; anions, with oxidation numbers for calcium and fluorine, +2, and &#8722;1.
Covalent Compounds
Covalent compounds are more challenging in the use of formalism. Water is a covalent compound consisting of two H atoms bonded to an O atom via polar covalent O&#8722;H bonds. The shared electrons making an O&#8722;H bond are more strongly attracted to the more electronegative O atom. So oxygen acquires a partial negative charge, compared to an O atom in elemental oxygen. As a result, H atoms in a water molecule exhibit a partial positive charge compared to hydrogen atoms in hydrogen gas. The sum of the partial negative and positive charges for each water molecule is zero, making the water molecule neutral.
If the polarization of shared electrons within the O&#8722;H bonds of water were complete&#8212;the result would be the complete transfer of electrons from H to O, and water would be an ionic compound composed of O2&#8722; anions and H+ cations. And so, the oxidations numbers for oxygen and hydrogen in water are &#8722;2 and +1, respectively. Applying this same logic to carbon tetrachloride, CCl4, yields oxidation numbers of +4 for carbon and &#8722;1 for chlorine. In the nitrate ion, NO3&#8722;, the oxidation number for nitrogen is +5 and that for oxygen is &#8722;2, summing to equal the &#8722;1 charge on the molecule:
<img alt="Static equilibrium diagram, &#931;Fx=0, showing forces on beam with supports, physics concept." src="/files/ftp_upload/11432/11432_Equation_1.png" style="text-align: center; height: 70px;" data-altupdatedat="1754904545000" data-alt="Mendelian inheritance, Punnett square, genetic diagram showing alleles and offspring trait prediction." />
Balancing Redox Equations
The unbalanced equation shown below describes the decomposition of sodium chloride:
<img alt="Chromatogram diagram showing amino acid separation using gas chromatography, peak analysis results." src="/files/ftp_upload/11432/11432_Equation_2.png" style="text-align: center; height: 48px;" data-altupdatedat="1754904545000" data-alt="Thermodynamics diagram of the Carnot cycle, PV graph, illustrating heat engine process stages." />
This reaction matches the criterion for a redox reaction, as the oxidation number for Na decreases from +1 to 0 (by undergoing reduction) and that for Cl increases from &#8722;1 to 0 (by undergoing oxidation). The equation case is easily balanced by adding the stoichiometric coefficient of 2 for the NaCl and Na:
<img alt="Dynamic equilibrium, &#931;F=0, diagram showing forces, lever balance, static analysis, physics concept." src="/files/ftp_upload/11432/11432_Equation_3.png" style="text-align: center; height: 48px;" data-altupdatedat="1754904545000" data-alt="BJT transistor symbol; diagram showing NPN transistor for electronics; key component in circuits." />
Redox reactions occurring in aqueous solutions are commonly encountered in electrochemistry, and many involve water or its ions, H+ (aq) and OH&#8722; (aq), either as reactants or products.
In these cases, equations representing redox reactions can be very challenging to balance merely by inspection, and the use of a systematic approach known as the half-reaction method is helpful. This approach involves the following steps:
<ol>
	<li>Split the equation into its component skeleton oxidation and reduction half-reactions.</li>
	<li>Balance each half-reaction for all elements other than O and H.</li>
	<li>Balance each half-reaction for O atoms by adding water molecules as needed in the equation.</li>
	<li>Balance each half-reaction for H atoms by adding protons as required for the equation.</li>
	<li>Finally, balance the charges on the elements by adding electrons as needed.</li>
	<li>Multiply the half-reactions by any integer needed to equalize the number of electrons lost in the oxidation half-reaction to the number of electrons gained in the reduction half-reaction.</li>
	<li>Add both half-reactions and simplify it further by canceling out common species on both sides of the equation.</li>
	<li>If the reaction occurs in an alkaline medium, add OH&#8722; ions to the equation obtained in step 7 to neutralize the protons (added in equal numbers to both sides of the equation) and simplify.</li>
	<li>Check the equation to ensure that the charges on the atoms are balanced.</li>
</ol>
This text is adapted from <a href="https://openstax.org/books/chemistry-2e/pages/17-1-review-of-redox-chemistry">Openstax, Chemistry 2e, Chapter 17: Introduction,</a> and <a href="https://openstax.org/books/chemistry-2e/pages/17-1-review-of-redox-chemistry">Openstax, Chemistry 2e, Section17.2: Review of Redox Chemistry</a>.

Electrochemistry is the science involved in the interconversion of electrical and chemical reactions. Such reactions are called reduction-oxidation, or ...

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KEY TERMS AND CONCEPTS

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 charge. Sodium chloride has Na at +1 and Cl at -1, and calcium fluoride has Ca at +2 and F at -1.

Covalent compounds need a closer look because electrons are shared. In water, the O-H bonds are polar, so oxygen pulls shared electrons more strongly than hydrogen does. That gives oxygen an oxidation number of -2 and hydrogen +1. The same idea gives carbon +4 and chlorine -1 in carbon tetrachloride, and nitrogen +5 with oxygen -2 in the nitrate ion, which adds up to the ion's -1 charge.

A simple redox example is the decomposition of sodium chloride. Sodium changes from +1 to 0, so it is reduced. Chlorine changes from -1 to 0, so it is oxidized. The equation is balanced by placing a coefficient of 2 in front of NaCl and Na.

Redox reactions in aqueous solution are often harder to balance because water, H+ ions, or OH- ions can appear as reactants or products. A systematic half-reaction method helps. First, split the equation into oxidation and reduction half-reactions. Then balance all elements except O and H, add water for oxygen, add protons for hydrogen, and add electrons to balance charge. After that, multiply the half-reactions so the electrons lost and gained are equal, add them together, and cancel species that appear on both sides.

If the reaction occurs in an alkaline medium, add OH- ions to neutralize the protons and simplify the equation. Finish by checking that both atoms and charge are balanced. This method gives a clear path for balancing redox equations in electrochemistry.

Redox Reactions and Half-Reaction Balancing

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

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...

key terms and concepts

Learn how oxidation numbers, redox changes, and the half-reaction method help balance redox equations in aqueous and alkaline solutions.

Redox Reactions and Half-Reaction Balancing | JoVE K12

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Balancing Chemical Equations and Reaction Outcomes

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Phase Transitions: Sublimation and Deposition

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Physical Properties of Matter

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Molecular Comparison of Gases, Liquids, and Solids

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Bonding in Metals

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Construct an explanation to classify the type of chemical bond that occurs (covalent, ionic, or metallic) in a given substance.

PS.PSc.2.1

Nomenclature of Alkanes

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Nomenclature of Alkenes

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Nomenclature of Aromatic Compounds with Multiple Substituents

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Use models to apply International Union of Pure and Applied Chemistry (IUPAC) conventions to name and write formulas for simple compounds.

PS.PSc.2.2

Understand types, properties, and structure of matter.

Redox Reactions and Half-Reaction Balancing

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