[Chemistry Class Notes] on Redox Reactions Pdf for Exam

The term redox is made up of two words: reduction and oxidation. A type of chemical reaction that involves a transfer of electrons between two species is called redox reaction. In these types of reactions, oxidation and reduction both take place together. 

Examples of Redox Reactions 

• Cellular Respiration – In respiration reaction, glucose reacts with oxygen and forms carbon dioxide and water, and releases energy that is stored in the cells. Glucose gets oxidised into carbon dioxide by losing hydrogens, while oxygen gets reduced to water by gaining hydrogens. The reaction is given below:

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• Combustion – Combustion is an exothermic redox chemical reaction, which occurs at high temperatures and in presence of an oxidant. Methane burns in presence of atmospheric oxygen and gives carbon dioxide and water energy. A well-explained combustion reaction of methane is given below:

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• Photosynthesis – Photosynthesis is a process by which plants prepare their food. In this reaction, carbon dioxide reacts with water in the presence of sunlight and chlorophyll and gives carbohydrates and oxygen. A well-explained reaction of photosynthesis is given below:

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• Corrosion – The rusting process is an example of corrosion. In this reaction, iron reacts with atmospheric oxygen in presence of moisture and forms iron oxide, which is also called rust. A well-explained reaction of rusting is given below:

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• Formation of Sodium Chloride – In this reaction, sodium reacts with chlorine and forms sodium chloride. In the reaction, the oxidation number of sodium increases from 0 to +1, while the oxidation number of chlorine atoms decreases from 0 to -1. It can also be explained in terms of gain and loss of electrons. A well-explained reaction is given below:

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What is Oxidation? 

Oxidation in terms of Oxygen transfer – Oxidation is the gain of oxygen. 

 

Example – 2Mg_(s) + O_2(g) ? 2MgO_(s) 

 

Oxidation in terms of electron transfer – Oxidation is the loss of electrons. In the above example, magnesium is losing two electrons and getting oxidised to form magnesium oxide. 

 

Oxidation in terms of Hydrogen transfer – Oxidation can be defined as loss of hydrogen.

 

Example – ()

 

Oxidation in terms of Oxidation number – Oxidation is an increase in the oxidation state or oxidation number of an atom in a reaction. Oxidation number defines the degree of oxidation of an atom in a chemical compound. 

 

Example –  ()

 

In the above example, the oxidation state of sodium is increasing from 0 to +1. Thus, oxidation is taking place and sodium is getting oxidised. }

What is Reduction?

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Reduction in terms of Oxygen transfer – Reduction is loss of oxygen. 

 

Example – ()

 

Reduction in terms of Electron transfer – Reduction is the gain of electrons. 

 

Example – 2Mg_(s) + O_2(g) ? 2MgO_(s)

 

In the above example, each atom of oxygen gains two electrons and forms two O^-2 anions. Thus, reduction takes place and oxygen gets reduced.  

 

Reduction in terms of hydrogen transfer – Reduction can be defined as the gain of hydrogen.

 

Example – ()

 

Reduction in terms of Oxidation number – Reduction is a decrease in the oxidation state or oxidation number of an atom in a reaction. 

 

Example – ()

 

In the above example, the oxidation state of chlorine is decreasing from 0 to -1. Thus, the reduction is taking place and chlorine is getting reduced. 

What is an Oxidising Agent and Reducing Agent?

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A substance that can oxidise another substance is called an oxidising agent. It is also called an oxidant. It oxidises another substance by accepting its electrons. Thus, we can say an oxidising agent is an electron acceptor group. An oxidising agent always reduces itself and oxidises another substance. Oxygen, hydrogen peroxide, and halogens are examples of some common oxidising agents. 

 

The oxidising agent can be defined as those groups which transfer oxygen atoms to the substrate, although it’s not true in every case, as in many redox reactions, oxidation-reduction takes place in absence of oxygen atoms as well. These oxidising agents are also called oxygenation reagents or oxygen–atom transfer (OAT) agents. MnO₄–, CrO₄-2, etc. are examples of these types of oxidising agents. You can notice here that these are all oxides. 

 

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A substance that can reduce another substance is called a reducing agent. It is also called reductant or reducer. It reduces another substance by donating its electrons. Thus, we can say that the reducing agent is an electron donor group. A reducing agent always oxidises itself and reduces another substance, like Lithium aluminium hydride (LiAlH₄), Hydrogen, sodium amalgam, etc.

 

Na(Hg)

 

Na(Hg) are examples of some common reducing agents.

 

The reducing agent can be defined as those groups which gain oxygen atoms from the substrate (or oxidising agent), although it is not true in every case, as in many redox reactions, oxidation-reduction takes place in absence of oxygen atoms. Examples of these reducing agents include formic acid, oxalic acid, sulfites, etc.

Oxidising Agent and Reducing Agent in a Redox Reaction 

Let’s understand oxidising and reducing agents by taking an example of a redox reaction. When aluminium reacts with iron(III) oxide in presence of heat, it gives aluminium oxide and molten iron metal. It is a redox reaction. The reaction is given below:

 

2Al_(s) + Fe₂O_3(s) ? Al₂O_3(s) + 2Fe_(I)

 

If you calculate the oxidation number for Al, then you see that it’s increasing from 0 to +3, which means that oxidation is taking place. Now, if you calculate it for iron, then you see that it’s decreasing from +3 to 0, which means a reduction is taking place. It is represented below in the reaction:

 

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You can see that Al is reducing Fe₂O₃ by removing its oxygen atoms. So, it means aluminium is a reducing agent, while Fe₂O₃ is oxidising Al by giving oxygen to it. So, it means Fe₂O₃ is acting as an oxidising agent. In redox reactions, reducing agents always convert into their conjugating oxidising agents in oxidation–reducing reaction. Thus, the products of this reaction will include a new oxidising agent and a new reducing agent. It is represented below in the reaction:

 

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As we know, the above reaction proceeds in the forward direction, so Al is a stronger reducing agent and Fe₂O₃ is a stronger oxidising agent than iron and Al₂O₃, respectively.

Another example of a redox reaction, in which oxidising agent and reducing agent have been indicated clearly, is as follows:

 

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What is the Oxidation Number? 

In order to keep track of electron shifts in chemical reactions involving the formation of covalent compounds, a more practical method of using oxidation numbers has been developed. Oxidation number can be defined as an imaginary or apparent charge developed over an atom of an element when it goes from its elemental free state to a combined state in molecules. 

 

Oxidation number denotes the oxidation state of an element in a compound, ascertained according to a set of rules formulated on the basis that the electron pair in a covalent bond belongs entirely to a more electronegative element. 

Rules for Calculating Oxidation Number

The following rules must be kept in mind while assigning the oxidation number to the atom in a molecule or in an ionic state: 

• Different elements in their elementary state or in allotropic form bear 0 as their oxidation number. For example, Nitrogen(N₂) and hydrogen(H₂), in their elemental state, will have zero oxidation state. Another example is, in diamond, graphite and buckminsterfullerene carbon have oxidation numbers ‘0’.

• The charge present on any monatomic ion is its oxidation number. For example, the oxidation number of Mg²⁺ is 2, whereas the oxidation number of Al³⁺ is 3.

• Hydrogen shows two oxidation states. When it combines with non-metals, it shows the oxidation state of +1, and when combined with metals and metal hydrides, it shows the oxidation state of -1. For example, in H₂S, the oxidation state of hydrogen is +1 because S is non-metal, whereas in MgH₂, the oxidation state of hydrogen is -1 as Mg is a metal.

• The most common oxidation state of oxygen is -2. When it combines with metals, it shows the oxidation state of -2, although it shows other oxidation states as well, such as in peroxide, it shows -1 oxidation state. With fluorine in OF₂, it shows the oxidation state of +2, and in O₂F₂, it shows the oxidation state of +1.

• When different elements unite to form compounds, more electronegative elements will show a negative oxidation state, whereas fewer electronegative elements will show a positive oxidation state. For example, in NaCl, Na shows a +1 oxidation state, whereas Cl shows a -1 oxidation state. Also, the oxidation number of Cl is -1 in HCl, but the oxidation number of Cl is +1 in HOCl. 

• The sum of oxidation numbers of all the atoms in a compound is zero. For example, in NaCl, the sum of oxidation states of Na and Cl is 0, which is +1-1=0.

• In complex ions, the sum of the oxidation numbers of all the atoms in the ionic state is equal to the charge on the ion. For example, in hexacyanoferrate (II) ion.

• Fe(CN)₆

• In Fe(CN)₆^-4, the sum of the oxidation state of Fe and CN is -4. In other words, in Radicals/Polyatomic ions, the oxidation state is the charge on the ion, which in turn, is equal to the algebraic sum of the charge on all the atoms. For example, in NH₄⁺, the sum of oxidation states of N and H is +1.

Apply all the above rules to the compound (NH₄)₂SO₄ to calculate the oxidation number of Nitrogen.

 

Let us suppose the oxidation state of nitrogen to be x in (NH₄)₂SO₄. 

 

As we know, the oxidation state of hydrogen is +1 and that of sulphate is -2, and the given compound carries zero charge. 

 

Therefore, the sum of all the oxidation states of elements will be zero.

 

2x + 2(+1×4) + (-2) = 2x + 6

 

2x + 6 = 0

 

x = -3

 

Hence, the oxidation number of Nitrogen in (NH₄)₂SO₄ is -3.

Types of Redox Reactions 

Redox reactions are of the following four types: 

Combination Reactions – In combination reactions, two or more molecules are combined together chemically to form a new substance (compound). For example, when we burn magnesium ribbon (or magnesium), it gives grey-black ash of magnesium oxide. Another example is the reaction between magnesium and nitrogen. 

E.g., 4 Fe+ 3O₂→2Fe₂O

Example –  ()

Decomposition Reaction – In a decomposition reaction, molecules or compounds break down into two or more than two simpler, chemically new substances. Combination and decomposition reactions are opposite of each other. For example, electrolysis of water. In the electrolysis of water, water breaks down into hydrogen and oxygen, which show completely different properties than water. 

Examples – 2H₂O electricity⟶ 2H₂ + O₂

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Displacement Reaction – In these reactions, more reactive metal displaces less reactive metal from its salt. In these reactions, products can be determined through reactivity series. Reactivity series is a series in which elements are arranged in decreasing order of their reactivity. It means the elements present at the top of this reactivity series are more reactive than the elements present at the bottom. 

The reaction of potassium with magnesium chloride is an example of a single displacement reaction. In this reaction, potassium displaces magnesium from its salt because potassium is more reactive than magnesium. Potassium is present at the top of the reactivity series and is the most reactive element.

Reaction – 2K + MgCl₂ ? 2KCl + Mg

Disproportionation Reactions – These are special types of redox reactions. In these reactions, an element in one oxidation state is simultaneously oxidised and reduced. In these reactions, one of the reacting substances always contains an element that can exist in at least three oxidation states. The decomposition of hydrogen peroxide is a common example of this type of reaction. The equation is given below:

E.g., 2 CuBr ? CuBr₂ + Cu

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Balancing of Redox Reactions – For balancing the redox reactions, the oxidation number method and half-reaction method are used.