What is the Science Behind Oxidation and Reduction?

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Oxidation and reduction are two of the most common chemical reactions. They are unbreakably linked, and many natural processes like combustion, fruit browning, and photosynthesis involve redox reactions.

Oxidation involves electron-stealing, while reduction involves electron-gaining. The amount of electrons stolen or gained defines an atom’s oxidation state.


Rusting is the oxidisation of iron and iron alloys (like steel). This causes the formation of a red-orange coloured substance called rust which weakens and eventually breaks down metals.

Often, we see rust on molds, dies and machinery – but what exactly is this chemical reaction that forms when metals come into contact with oxygen or moisture? This is an important concept to know because it can help you prevent rust from forming in the first place and keep your equipment, cars and buildings safe.

It is a common process that occurs naturally when the iron in metal comes into contact with water and oxygen. The resulting reaction transforms the iron into a compound known as ferric oxide – rust.

The oxidation process happens because the iron in the metal is attracted to the oxygen molecules. Oxygen has a positive charge and the iron has a negative charge, so when they combine, they create a chemical reaction that produces ferric oxide.

This is a very destructive process because it weakens and crumbles metal and can make it brittle and breakable. This can be a serious issue for industries that use valuable metals in their operations, as it could cause them to become damaged and unusable over time.

There are a number of different types of corrosion, but the most common is rusting. Some other types include pitting and cavity corrosion, contact corrosion and crevice corrosion.

For instance, pitting and cavity corrosion occur when iron or stainless steel gets into contact with another corroding metal or material. These are usually in confined spaces, like the space between two nuts or bolts. They can progressively destroy the integrity of the metal and lead to its collapse or even detachment from its support.

Corrosion can also be a dangerous process for humans because it exposes them to toxic fumes and vapors, which can lead to severe health problems if the exposure is prolonged. For this reason, it is very important to take precautions to avoid exposing people to rust or other types of corrosion.

In order to protect their valuable equipment from rusting, it is essential that mold makers and plastic molding companies have good hygiene practices in place. This includes using a rust-resistant coating on their machines and cleaning them regularly to remove any remaining moisture that could be contributing to the rusting process.

Oxidising Agents

Oxidising agents are substances that accept electrons from another chemical species in a reaction and give oxygen to it or remove hydrogen from it. They are used in a variety of reactions, including the burning of different types of fuels, electrochemical processes for the extraction of metals and non-metals from waste products, battery operation, corrosion of materials, etc.

Typically, oxidising agents are elements that have a high electron affinity. These include iodine, bromine, chlorine, fluorine and the halogens (group 17 on the periodic table). The number of valence electrons also has an effect on an oxidising agent’s power to gain electrons, which is why the halogens have the highest electronegativities.

An oxidising agent can be inorganic or organic. The latter includes organic peroxides such as dibenzoyl peroxide and methyl ethyl ketone peroxide, widely used as catalysts in the manufacture of plastics.

However, oxidising agents have the potential to become explosive. As a result, they are classified as hazardous and require careful handling and storage.

Most oxidising agents are exothermic, meaning that they produce heat quickly. This makes them a serious fire hazard.

They can be found in many different forms, including anhydrous, mono, tri and tetrahydrate. One commonly used oxidising agent is sodium perborate. It can be purchased in powder or patented commercial combinations.

The oxidising powers of an oxidising agent can vary depending on the type of substance being oxidised, and how much it is able to bind oxygen. For example, sodium chlorate is an extremely powerful oxidising agent, but calcium carbonate and calcium sulphate are not as strong.

An oxidising agent’s ability to oxidise is also influenced by its standard reduction potential, which is the difference between its maximum possible oxidation state and its minimum reduction potential. The higher the standard reduction potential, the more it can oxidise.

In some cases, a strong oxidising agent can react vigorously with other compounds, generating heat and possibly gaseous products that can pressurize a closed container. This is an important characteristic of oxidising agents, as it means that they may be able to contribute to other chemical reactions.

Reduction Agents

Reduction agents are substances that have the ability to donate electrons to another substance. These substances are often used in redox reactions (also known as electron transfer reactions).

Reducing agents tend to be atoms or molecules with a relatively low oxidation state and a high electronegativity, which means they can easily lose or donate their outer electrons. These reducing agents can be found in many different forms, from compounds like zinc oxide and iron (Fe) to hydrogen gas and magnesium oxide.

Species that have a relatively long atomic radius, such as the alkali and alkaline earth metals in Groups 1a and 2a of the periodic table, are also strong reducing agents because they have a long distance from their nucleus to their valence electrons. This allows them to ‘lose’ their electrons easily, so that they become a positively charged metal ion with a higher oxidation number than the oxygen that is in contact with them.

A species’ reducing power is usually assessed using its standard reduction potential, which is shown in Table 11 below. Species that lie above H2 are stronger reducing agents than H2, and species that lie below it are more oxidizing agents than H2.

These reduction potentials are measured in volts or millivolts at 298 K under 1 atm using 1 molar solution. In general, the more negative a substance’s reduction potential is, the stronger it will be as a reducing agent.

Although the standard reduction potentials are only a guide to the oxidizing and reducing power of certain substances, they do provide a useful ranking. For example, sodium is a strong reducing agent but is weak as an oxidizing agent because it has a large oxidation potential.

Other elements and compounds that have a high oxidation potential can be considered reducing agents as well, because they have the ability to lose their outer electrons easily. These reducing agents include hydrides, which are compounds that contain hydrogen in the formal -1 oxidation state.

Other reducing agents include some of the compounds that are found in proteins, such as DTT, glutathione and TCEP. These reducing agents can help to cleave the disulfide bond between cysteine amino acids, which is an important step in protein synthesis.

Redox Reactions

Oxidation and reduction are the two types of chemical reactions that take place in redox (oxidation-reduction) reactions. They are important in chemistry and biology, but also play a role in our lives in many ways.

Redox reactions are the basis for much of the energy that we use and produce in our bodies. Cellular respiration, for instance, uses redox reactions to break down glucose into ATP, our body’s primary source of energy. Redox reactions also enable the production of caustic soda and chlorine, which are essential for sanitizing water and bleaching materials.

In a redox reaction, one substance is oxidized (lost electrons) and another is reduced (gains electrons). Both oxidation and reduction occur simultaneously in redox reactions and cannot happen independently.

This makes redox reactions difficult to calculate, so we often use half-reactions to help us understand the oxidation and reduction of a compound. These half-reactions are balanced so that the net change is cancelled out.

However, redox reactions are also complex, and it isn’t uncommon for some of them to be difficult to understand. For this reason, we sometimes try to simplify them as much as possible.

The basic idea behind redox is that the atoms or molecules involved in a reaction have to be electron-thirsty, and that they need an atom or molecule to transfer electrons to. Alternatively, they must have an atom or molecule that wants to strip electrons away.

We can see this in practice if we look at the reaction between carbon dioxide and oxygen. The carbon atoms lose electrons, while the oxygen atoms gain electrons. This is because oxygen is much more electronegative than carbon. Its electrons “hog” the bonds between carbon atoms, and those bonds become weaker.

It is also possible for both carbon and oxygen to bind together, such as in the displaystyle textOC=OC bonds of water. The hydrogen atoms gain electrons from the oxygen atoms in these bonds, so they are stronger and more stable than before.

Redox reactions are significant in geology, too, as they are associated with mineral deposition and mobilization. Uranium deposits, Moqui marbles and other minerals are famous examples of redox reactions that have shaped the Earth’s surface.

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