Redox Reactions
Definition of redox: Redox reactions are types of chemical reactions where two very important things happen at the same time: oxidation, where a substance loses electrons, and reduction, where a different substance gains those electrons.
Electron transfer: These reactions always involve the movement, or transfer, of electrons from one chemical substance to another. This flow of electrons is what makes redox reactions so important in chemistry and everyday life.
Defining Oxidation
Multiple definitions: Oxidation can be explained in several different ways depending on the situation or reaction.
Addition of oxygen: One way to define oxidation is when a substance gains oxygen atoms. For example, when magnesium burns in air, it reacts with oxygen to form a new substance called magnesium oxide.
Removal of hydrogen: Oxidation can also mean that a substance loses hydrogen atoms. An example is when hydrogen sulfide loses hydrogen and changes into sulfur.
Loss of electrons: Another way to understand oxidation is that a substance loses electrons during a reaction. For instance, when a potassium atom loses an electron, it becomes a potassium ion.
Increase in oxidation number: If the oxidation number (a way to track electrons) of an atom increases during a reaction, it means oxidation has occurred. For example, when magnesium reacts with sulfuric acid, its oxidation number goes up.
Defining Reduction
Opposite of oxidation: Reduction is the exact opposite of oxidation. It happens when a substance gains electrons instead of losing them.
Removal of oxygen: Reduction can involve removing oxygen atoms from a substance. For example, copper(II) oxide can be reduced to copper metal by taking away its oxygen.
Addition of hydrogen: Reduction can also happen when a substance gains hydrogen atoms. For instance, chlorine can combine with hydrogen to form hydrogen chloride.
Gain of electrons: During reduction, a substance gains electrons. An example is when a chlorine molecule gains electrons and becomes two chloride ions.
Decrease in oxidation number: If the oxidation number of a substance decreases during a reaction, it means that reduction has occurred. For example, copper(II) ions are reduced by magnesium to form copper metal.
Oxidising and Reducing Agents
Oxidising Agents: An oxidising agent is a substance that helps another substance to get oxidised. It does this by accepting electrons from the other substance. In simpler words, the oxidising agent takes electrons away from something else, which causes that something else to lose electrons and undergo oxidation.
Oxidising Agent Examples: Some common examples of oxidising agents include oxygen gas (O₂), chlorine gas (Cl₂), and potassium manganate(VII) (KMnO₄). These substances are very good at accepting electrons from other materials during chemical reactions.
Oxidising Agent Role Examples: For instance, oxygen acts as an oxidising agent when it reacts with burning magnesium. The oxygen accepts electrons from the magnesium, allowing magnesium oxide to form. Another example is chlorine gas, which acts as an oxidising agent when it reacts with hydrogen sulfide, accepting electrons and leading to the formation of new products.
Reducing Agents: A reducing agent is a substance that causes another substance to get reduced. It does this by donating electrons to the other substance. In other words, the reducing agent gives away electrons, helping the other substance gain electrons and undergo reduction.
Reducing Agent Examples: Some typical examples of reducing agents are hydrogen gas (H₂), magnesium metal (Mg), and zinc metal (Zn). These substances are very good at donating their electrons to other substances during reactions.
Reducing Agent Role Examples: For example, hydrogen gas can reduce copper(II) oxide (CuO) to copper metal by giving it electrons. Magnesium metal can reduce copper(II) sulfate (CuSO₄) to produce copper metal by a similar process.
Half-Equations
Representation of Redox: Redox reactions involve both oxidation (loss of electrons) and reduction (gain of electrons). These reactions can be broken down into two simpler parts, called half-equations. Each half-equation shows either the oxidation or the reduction part of the overall reaction.
Oxidation Half-Equations: An oxidation half-equation shows how a substance loses electrons during a reaction. For example:
- Magnesium atom losing electrons: Mg → Mg²⁺ + 2e⁻
- Bromide ions losing electrons to form bromine gas: 2Br⁻ → Br₂ + 2e⁻
Reduction Half-Equations: A reduction half-equation shows how a substance gains electrons during a reaction. For example:
- Copper ions gaining electrons to form copper metal: Cu²⁺ + 2e⁻ → Cu
- Chlorine gas gaining electrons to form chloride ions: Cl₂ + 2e⁻ → 2Cl⁻
Oxidation Numbers
Tracking Electrons: Oxidation numbers are useful tools that scientists use to keep track of how many electrons an atom gains or loses during a chemical reaction. They help explain which atoms are oxidised and which are reduced.
Elemental State: When elements like magnesium metal or oxygen gas are in their pure form (not combined with anything else), their oxidation number is always zero.
Monatomic Ions: For simple ions made of just one atom, the oxidation number is the same as the ion’s charge. For example, a sodium ion (Na⁺) has an oxidation number of +1.
Oxygen Rule: Oxygen almost always has an oxidation number of -2 when it is part of a compound. However, in special compounds called peroxides (like hydrogen peroxide), oxygen has an oxidation number of -1 instead.
Hydrogen Rule: Hydrogen usually has an oxidation number of +1 when it forms compounds with non-metals. But in special compounds called metal hydrides, hydrogen has an oxidation number of -1.
Neutral Compound Sum: In a neutral compound (a compound with no overall charge), the sum of all the oxidation numbers of the atoms must add up to zero.
Polyatomic Ions: In ions made up of more than one atom (polyatomic ions), the sum of the oxidation numbers of all the atoms must equal the overall charge of the ion.
Group 1 and Group 2 Elements: Elements from Group 1 of the periodic table, like sodium and potassium, always have an oxidation number of +1 in their compounds. Elements from Group 2, like magnesium and calcium, always have an oxidation number of +2.
Halogens: Halogens, such as fluorine, chlorine, and bromine, usually have an oxidation number of -1 when they form compounds with other elements.
Oxidation Changes: If a substance’s oxidation number increases during a reaction, it means that oxidation has taken place (the substance has lost electrons).
Reduction Changes: If a substance’s oxidation number decreases during a reaction, it means that reduction has happened (the substance has gained electrons).
Spontaneous Redox Reactions and Electrochemical Series
Spontaneity: Some redox reactions happen naturally, without the need for extra help like heating or electricity. Whether a reaction happens on its own depends mainly on the types of metals involved.
Electrochemical Series: Scientists have made a special list called the electrochemical series. This list arranges metals based on how easily they lose electrons.
Top Series Metals: Metals that are found at the top of the electrochemical series, such as potassium and sodium, lose electrons very easily and are strong reducing agents.
Bottom Series Metals: Metals found near the bottom of the electrochemical series, like gold and platinum, do not lose electrons easily. This makes them weak reducing agents.
Metal Displacement: A metal that loses electrons more easily (is more electropositive) can displace a less electropositive metal from a solution during a chemical reaction.
Displacement Example: For example, zinc metal can push copper metal out of a solution of copper(II) sulfate because zinc is better at giving away electrons compared to copper.
Examples of Redox Reactions
Metal Displacement: Magnesium reacts with copper(II) ions in a solution. During this reaction, magnesium gives away electrons to form magnesium ions, while copper(II) ions gain electrons and turn into copper metal.
Reaction with Oxygen: When magnesium burns in oxygen gas, it forms magnesium oxide. In this reaction, magnesium atoms lose electrons to oxygen atoms, clearly showing oxidation.
Reaction with Halogens: When chlorine gas reacts with hydrogen sulfide, it produces sulfur and hydrogen chloride gas. In this reaction, chlorine gets reduced (gains electrons) and hydrogen sulfide gets oxidised (loses electrons).
Electrochemical Cell Reactions: In devices like the Daniel cell, redox reactions happen separately at two different electrodes. One electrode carries out oxidation, and the other carries out reduction. The movement of electrons from one electrode to the other generates electricity.
Important Note
Not all reactions are redox: Not every chemical reaction is a redox reaction. If there is no transfer of electrons or no change in oxidation numbers, such as in neutralization (acid-base reactions) or precipitation (formation of solids from solutions), then it is not considered a redox reaction.