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

K12 Standards

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NC K-12 Science Standard Course of Study

9-12

Physical Science

Matter and its Interactions

PS.PSc.2.3

Balancing Redox Equations

Cómo proporciona una batería electricidad para operar dispositivos portátiles? Por qué un aguacate una vez cortado se vuelve marrón o el metal se oxida bajo el agua? Estos procesos son impulsados por tipos específicos de reacciones químicas que involucran la transferencia de electrones de un átomo a otro.El átomo que pierde electrones se oxida, mientras que el átomo que gana electrones se reduce. Estas reacciones se denominan reacciones de oxidación-reducción o redox y se caracterizan por cambios en el estado de oxidación de uno o más reactivos. La interacción entre el magnesio y el ácido clorhídrico concentrado es un ejemplo de reacción redox.Aquí, el magnesio se oxida a iones de magnesio 2, y los protones se reducen a gas hidrógeno. Las ecuaciones químicas simples se pueden equilibrar fácilmente. Sin embargo, como la mayoría de las ecuaciones redox son complejas, se debe justificar el número de electrones perdidos y ganados.Al equilibrar ecuaciones redox, es importante observar la ley de conservación de la masa. La cantidad de cada elemento y cualquier ganancia o pérdida de electrones deben equilibrarse en ambos lados de la reacción. Por lo tanto, si se reduce un reactivo, otro debe oxidarse.El método de semirreacciones se utiliza para equilibrar con éxito las ecuaciones redox. Considere la reacción del permanganato con oxalato. Para soluciones acuosas ácidas, asigne los estados de oxidación y divida la reacción en sus semirreacciones componentes.Luego, equilibre cada media reacción omitiendo el hidrógeno y el oxígeno. Aquí, el manganeso ya está equilibrado, pero el oxalato se necesita un coeficiente de dos;en segundo lugar, equilibre los átomos de oxígeno con la adición de moléculas de agua;y en tercer lugar, equilibre los átomos de hidrógeno añadiendo protones cuando sea necesario. Equilibre las cargas agregando electrones.Dado que el manganeso se reduce, se agregan cinco electrones en el lado del reactivo. Por el contrario, el oxalato se oxida;por lo tanto, se añaden dos electrones en el lado del producto. Multiplique las semirreacciones por un número entero para que el número de electrones agregados sea igual.Por último, agregue y simplifique las semirreacciones balanceadas cancelando especies en ambos lados para obtener una reacción redox balanceada. Para soluciones acuosas básicas, el procedimiento es similar pero incluye un paso adicional. Aquí, se agrega un número igual de iones hidroxilo a ambos lados de la reacción para neutralizar los protones antes de sumar las semirreacciones balanceadas.

balanceo de ecuaciones redox

La electroquímica es la ciencia involucrada en la interconversión de reacciones eléctricas y químicas. Tales reacciones se llaman de reducción-oxidación, o reacciones redox. Estas importantes reacciones se definen por cambios en los estados de oxidación para uno o más elementos reactantes e incluyen un subconjunto de reacciones que implican la transferencia de electrones entre especies reactantes. La electroquímica como campo ha evolucionado para ofrecer suficientes conocimientos sobre los principios fundamentales de la química redox y múltiples tecnologías que van desde procesos metalúrgicos a escala industrial hasta baterías recargables y robustas para vehículos eléctricos. Como las reacciones que involucran la transferencia de electrones son esenciales para el estudio de la electroquímica, una breve revisión de la química redox incluye lo siguiente.
 Números de oxidación
 Por definición, una reacción redox implica un cambio en el número de oxidación o en el estado de oxidación de uno o varios reactivos. El número de oxidación de un elemento es una evaluación de cómo el entorno electrónico de sus átomos difiere en comparación con los átomos del elemento puro. Por esta definición, un átomo en un elemento lleva un número de oxidación de cero. Para un átomo, el número de oxidación es igual a la carga del átomo en el compuesto si el compuesto era iónico. Así, la suma de los números de oxidación para todos los átomos de una molécula es igual a la carga de la molécula.
 Compuestos iónicos
 Los compuestos iónicos simples son los ejemplos más sencillos de este formalismo ya que los elementos tienen números de oxidación iguales a sus cargas iónicas. El cloruro de sodio, NaCl, está compuesto por cationes Na+ y aniones Cl−, siendo los números de oxidación para el sodio y el cloro +1 y −1, respectivamente. El fluoruro de calcio, CaF2, se compone de cationes Ca2+ cationes y aniones F−, con números de oxidación para calcio y flúor, +2, y −1 respectivamente.
 Compuestos covalentes
 Los compuestos covalentes son más desafiantes en el uso del formalismo. El agua es un compuesto covalente que consta de dos átomos de H unidos a un átomo de O a través de enlaces covalentes polares O−H. Los electrones compartidos que forman un enlace O−H son más atraídos al átomo O más electronegativo. Por lo tanto, el oxígeno adquiere una carga negativa parcial, en comparación con un átomo de O en el oxígeno elemental. Como resultado, los átomos de H de una molécula de agua presentan una carga positiva parcial en comparación con los átomos de hidrógeno del gas hidrógeno. La suma de las cargas parciales negativas y positivas para cada molécula de agua es cero, lo que hace que la molécula de agua sea neutra.
 Si la polarización de los electrones compartidos dentro de los enlaces O−H del agua estuviera completa, el resultado sería la transferencia completa de electrones de H a O, y el agua sería un compuesto iónico formado por aniones O2− y cationes H+. Y así, los números de oxidación para el oxígeno y el hidrógeno en agua son −2 y +1, respectivamente. Al aplicar esta misma lógica al tetracloruro de carbono, CCl4, se obtienen números de oxidación de +4 para el carbono y −1 para el cloro. En el ion nitrato, NO3−, el número de oxidación para el nitrógeno es +5 y el del oxígeno es −2, sumando para igualar la carga −1 sobre la molécula:
 <img alt="Eq1" src="/files/ftp_upload/11432/11432_Equation_1.png" style="text-align: center; height: 70px;" />
 Balanceando ecuaciones redox
 La ecuación no balanceada que se muestra a continuación describe la descomposición del cloruro de sodio:
 <img alt="Eq2" src="/files/ftp_upload/11432/11432_Equation_2.png" style="text-align: center; height: 48px;" />
 Esta reacción coincide con el criterio de una reacción redox, ya que el número de oxidación para el Na disminuye de +1 a 0 (por reducción) y el del Cl aumenta de −1 a 0 (por oxidación). El caso de la ecuación se balancea fácilmente añadiendo el coeficiente estequiométrico 2 para el NaCl y el Na:
 <img alt="Eq3" src="/files/ftp_upload/11432/11432_Equation_3.png" style="text-align: center; height: 48px;" />
 Las reacciones redox que ocurren en soluciones acuosas son comúnes en electroquímica, y muchas involucran agua o sus iones, H+ (ac) y OH− (ac), ya sea como reactivos o productos.
 En estos casos, las ecuaciones que representan reacciones redox pueden ser muy difíciles de balancear simplemente por inspección, y el uso de un enfoque sistemático conocido como método de la semirreacción es útil. Este enfoque implica los siguientes pasos: 
 <ol>
 <li>Divida la ecuación en las semirreacciones de oxidación y reducción del esqueleto que la compone.</li>
 <li>Balancee cada semirreacción para todos los elementos que no sean O y H.</li>
 <li>Balancee cada semirreacción para los átomos de O agregando moléculas de agua según sea necesario en la ecuación.</li>
 <li>Equilibre cada semirreacción para los átomos de H agregando protones según sea necesario para la ecuación.</li>
 <li>Finalmente, balancee las cargas en los elementos agregando electrones según sea necesario.</li>
 <li>Multiplique las semirreacciones por cualquier entero necesario para igualar el número de electrones perdidos en la semirreacción de oxidación al número de electrones ganados en la semirreacción de reducción.</li>
 <li>Sume ambas semirreacciones y simplifique aún más cancelando especies comunes en ambos lados de la ecuación.</li>
 <li>Si la reacción se produce en un medio alcalino, añada iones OH− a la ecuación obtenida en el paso 7 para neutralizar los protones (añadidos en números iguales a ambos lados de la ecuación) y simplifique.</li>
 <li>Compruebe la ecuación para asegurarse de que las cargas en los átomos están balanceadas.</li>
 </ol>
 Este texto está adaptado de <a href="https://openstax.org/books/chemistry-2e/pages/17-1-review-of-redox-chemistry"> Openstax, Química 2e, Capítulo 17: Introducción,</a> y <a href="https://openstax.org/books/chemistry-2e/pages/17-1-review-of-redox-chemistry"> Openstax, Química 2e, Section17.2: Revisión de la Química Redox</a>.

Balanceo de Ecuaciones Redox

La electroquímica es la ciencia involucrada en la interconversión de las reacciones eléctricas y químicas. Tales reacciones se denominan reducción-oxidación, o ...

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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

balancing redox equations

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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Construct an explanation to classify matter as a pure substance or mixture; homogeneous or heterogeneous; element or compound; solution, colloid or suspension.

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Clausius-Clapeyron Equation

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Phase Transitions: Melting and Freezing

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Phase Transitions: Sublimation and Deposition

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Heating and Cooling Curves

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Phase Diagrams

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Use models to compare the phases of matter and the physical changes they undergo.

PS.PSc.1.2

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Physical and Chemical Properties of Matter

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Physical Properties of Matter

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Molecular Comparison of Gases, Liquids, and Solids

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Intermolecular Forces

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Comparing Intermolecular Forces: Melting Point, Boiling Point, and Miscibility

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Carry out investigations to compare physical and chemical properties of matter.

PS.PSc.1.3

The Atomic Theory of Matter

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Subatomic Particles

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The Bohr Model

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Emission Spectra

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The de Broglie Wavelength

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The Uncertainty Principle

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Use models to interpret the data presented in Bohr diagrams and electron dot diagrams for neutral atoms of elements 1 through 18.

PS.PSc.1.4

Ions and Ionic Charges

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Use models to compare representations of atoms, ions, and isotopes.

PS.PSc.1.5

The Periodic Table

As early chemists discovered more elements, they realized that various elements could be grouped by their similar chemical behaviors. One such grouping ...

Periodic Classification of the Elements

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Atomic Radii and Effective Nuclear Charge

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Ionic Radii

Ionic radius is the measure used to describe the size of an ion. A cation always has fewer electrons and the same number of protons as the parent atom; it ...

Ionization Energy

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Electron Affinity

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Alkali Metals

Group 1 elements are soft and shiny metallic solids. They are malleable, ductile, and good conductors of heat and electricity. The melting points of the ...

Halogens

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Noble Gases


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Use the Periodic Table as a model to predict the relative properties (metallic/nonmetallic character, ionic charge, and reactivity) and arrangement of elements based on the pattern of valence electrons in the outermost energy levels of atoms.

PS.PSc.1.6

Types of Chemical Bonds

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Ionic Bonding and Electron Transfer

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Covalent Bonding and Lewis Structures

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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

In the late 19th-century, the number of new chemical compounds discovered increased tremendously. Hence, the necessity arose to develop a naming system ...

Nomenclature of Alkenes

The IUPAC naming system for alkenes replaces -an- with -en- in the corresponding parent alkanes. Accordingly, a simple alkene replaces the -ane suffix of ...

Nomenclature of Alkynes

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Structure and Nomenclature of Alcohols and Phenols

Overview 
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Structure and Nomenclature of Ethers

Structure and Bonding
Ethers are organic compounds with an ether functional group which is characterized by an oxygen atom connected to two &#8212; ...

IUPAC Nomenclature of Aldehydes

Aldehydes are named based on the systematic nomenclature rules set by the IUPAC. For acyclic aldehydes, the longest carbon chain containing the aldehydic ...

IUPAC Nomenclature of Ketones

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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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