Class 11 Chemistry Redox Reaction Notes Pdf Download

[Karina Edling]

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of reactions in which oxidation and reduction reactions occur simultaneously. Redox refers to an integral area of Chemistry. It is an essential reaction that concerns physical as well as biological phenomena.

The number of parts by mass of an element that displaces or reacts from a compound containing 8 parts by mass of oxygen, 1.008 parts by mass of hydrogen, and 35.5 parts by mass of chlorine is known as the equivalent weight of that element.

The actual number of hydrogen or hydroxide ion-exchanged in the reaction is the valence factor in an acid-base reaction. More hydrogen or hydroxide ions may be replaceable in the acid or base than what is actually replaced in the reaction. The number of hydrogen ions replaced by each molecule of the base from the acid is denoted by v. f.

According to the law, one equivalent of one element must combine with one equivalent of another. A chemical reaction results in the same number of equivalents or milliequivalents of products when equivalents or milliequivalents of reactants react in the same proportion.

Calcium and Magnesium bicarbonates cause temporary hardness while calcium and Magnesium chlorides and sulphates cause permanent hardness. The water sample can be softened by boiling or by using washing soda.

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If you have seen a piece of rusty metal then you have seen the end result of a redox reaction (iron and oxygen forming iron oxide). Redox reactions are also used in electrochemistry and in biological reactions.

When some reactions occur, an exchange of electrons takes place. It is this exchange of electrons that leads to the change in charge that we noted in grade 10 (chapter 18, reactions in aqueous solution). When an atom gains electrons it becomes more negative and when it loses electrons it becomes more positive.

Oxidation is the loss of electrons from an atom, while reduction is the gain of electrons by an atom. In a reaction these two processes occur together so that one element or compound gains electrons while the other element or compound loses electrons. This is why we call this a redox reaction. It is a short way of saying reduction-oxidation reaction!

Before we look at redox reactions we need to first learn how to tell if a reaction is a redox reaction. In grade 10 you learnt that a redox reaction involves a change in the charge on an atom. Now we will look at why this change in charge occurs.

By giving elements an oxidation number, it is possible to keep track of whether that element is losing or gaining electrons during a chemical reaction. The loss of electrons in one part of the reaction must be balanced by a gain of electrons in another part of the reaction.

This is an ionic compound composed of \(\textCa^2+\) and \(\textCl^-\) ions. Using rule 2 the oxidation number for the calcium ion is \(\text+2\) and for the chlorine ion it is \(-\text1\).

This is an ionic compound composed of \(\textMg^2+\) and \(\textSO_4^2-\) ions. Using rule 2 the oxidation number for the magnesium ion is \(\text+2\). In the polyatomic \(\textSO_4^2-\) ion, the sum of the oxidation numbers must be \(-\text2\) (rule \(\text4\)).

By looking at how the oxidation number of an element changes during a reaction, we can easily see whether that element is being oxidised (lost electrons) or reduced (gained electrons).

If the oxidation number of a species becomes more positive, the species has been oxidised and if the oxidation number of a species becomes more negative, the species has been reduced.

As a reactant, magnesium has an oxidation number of zero, but as part of the product magnesium chloride, the element has an oxidation number of \(\text+2\). Magnesium has lost two electrons and has therefore been oxidised (note how the oxidation number becomes more positive). This can be written as a half-reaction. The half-reaction for this change is:

As a reactant, chlorine has an oxidation number of zero, but as part of the product magnesium chloride, the element has an oxidation number of \(-\text1\). Each chlorine atom has gained an electron and the element has therefore been reduced (note how the oxidation number becomes more negative). The half-reaction for this change is:

You can remember this by thinking of the fact that when a compound is oxidised, it causes another compound to be reduced (the electrons have to go somewhere and they go to the compound being reduced).

A redox reaction is one involving oxidation and reduction, where there is always a change in the oxidation numbers of the elements involved. Redox reactions involve the transfer of electrons from one compound to another.

A recommended experiment for informal assessment on redox reactions is also included. This experiment is split into three parts. Each part looks at a different type of redox reaction (displacement, synthesis and decomposition). Learners looked at the second two reactions in grade 10.

When burning magnesium ribbon remind learners to not look directly at the flame. Also, since this experiment involves Bunsen burners, learners must work in a well ventilated room and take the usual care with loose scarves and long hair.

Hydrogen peroxide can cause serious chemical burns. Learners must work carefully with this substance. If they spill any on themselves then they must rinse the affected area with plenty of water and call you. If necessary they may need to go to the bathroom to rinse affected clothing off. You or a learner should accompany them.

\(\textCu^2+\) ions from the \(\textCuSO_4\) solution are reduced to form copper metal. This is what you saw on the zinc crystals. The reduction of the copper ions (in other words, their removal from the copper sulfate solution), also explains the change in colour of the solution (copper ions in solution are blue). The equation for this reaction is:

Using what you have learnt about oxidation numbers and redox reactions we can balance redox reactions in the same way that you have learnt to balance other reactions. The following worked examples will show you how.

The reaction for iron is the oxidation half-reaction as iron became more positive (lost electrons). The reaction for chlorine is the reduction half-reaction as chlorine has become more negative (gained electrons).

In the example above we did not need to know if the reaction was taking place in an acidic or basic medium (solution). However if there is hydrogen or oxygen in the reactants and not in the products (or if there is hydrogen or oxygen in the products but not in the reactants) then we need to know what medium the reaction is taking place in. This will help us to balance the redox reaction.

If a redox reaction takes place in an acidic medium then we can add water molecules to either side of the reaction equation to balance the number of oxygen atoms. We can also add hydrogen ions to balance the number of hydrogen atoms. We do this because we are writing the net ionic equation (showing only the ions involved and often only the ions containing the elements that change oxidation number) for redox reactions and not the net reaction equation (showing all the compounds that are involved in the reaction). If a Bronsted acid is dissolved in water then there will be free hydrogen ions.

If a redox reaction takes place in an basic medium then we can add water molecules to either side of the reaction equation to balance the number of oxygen atoms. We can also add hydroxide ions (\(\textOH^-\)) to balance the number of hydrogen atoms. We do this because we are writing the net ionic equation (showing only the ions involved and often only the ions containing the elements that change oxidation number) for redox reactions and not the net reaction equation (showing all the compounds that are involved in the reaction). If a Bronsted base is dissolved in water then there will be free hydroxide ions.

In the first reaction we have \(\text2\) chromium atoms and \(\text7\) oxygen atoms on the left hand side. On the right hand side we have \(\text1\) chromium atom and no oxygen atoms. Since we are in an acidic medium we can add water to the right hand side to balance the number of oxygen atoms. We also multiply the chromium by \(\text2\) on the right hand side to make the number of chromium atoms balance. \[\textCr_2\textO_7^2- \rightarrow 2\textCr^3+ + 7\textH_2\textO\]

Now we have hydrogen atoms on the right hand side, but not on the left hand side so we must add \(\text14\) hydrogen ions to the left hand side (we can do this because the reaction is in an acidic medium): \[\textCr_2\textO_7^2- + 14\textH^+ \rightarrow 2\textCr^3+ + 7\textH_2\textO\] We do not use water to balance the hydrogens as this will make the number of oxygen atoms unbalanced.

The reaction involving sulfur is the oxidation half-reaction as sulfur became more positive (lost electrons). The reaction for chromium is the reduction half-reaction as chromium has become more negative (gained electrons).

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