0620 · 6.4Redox

Particles and oxidation numbers

0:00In this reaction, oxygen moves from copper(II) oxide to hydrogen. One substance loses oxygen while another gains it. Those two changes belong to the same process, called a redox reaction. By the end of this video, you will recognise redox using oxygen, electrons and oxidation numbers, and explain what each substance does. We need two starting ideas about particles. An atom is electrically neutral because it has equal numbers of protons and electrons.

0:25Losing electrons produces a positive ion, while gaining them produces a negative ion. During chemical reactions, electrons can move between particles, but the total electrical charge must remain balanced. An oxidation number is a bookkeeping value assigned to an element in a substance. It helps us track changes during reactions, even when electrons are not shown in the equation. For a simple ion, this value matches its charge. In compounds, however, an oxidation number is not necessarily an actual charge on an atom.

0:54Roman numerals in compound names indicate the oxidation number of an element. Iron(II) oxide contains iron with oxidation number positive two. In iron(III) oxide, the corresponding value is positive three. The names use Roman numerals because iron can form compounds with different oxidation numbers, so the numeral distinguishes them. The numeral does not tell you how many atoms appear in the formula. Iron(III) oxide contains two iron atoms for every three oxygen atoms.

1:19Its name identifies the oxidation number of iron, not the number of iron atoms. Keep those two jobs separate when reading a name and a formula.

Redox through oxygen transfer

1:28At Core level, oxidation is defined as gain of oxygen, while reduction is loss of oxygen. A redox reaction involves simultaneous oxidation and reduction. Simultaneous means that both processes occur together as parts of the same reaction, rather than as two unrelated events. Our example is copper(II) oxide plus hydrogen gives copper plus water. The balanced symbol equation is on the board. Compare copper(II) oxide with the copper produced: oxygen has been removed, so the oxide is reduced.

1:58Hydrogen gains oxygen to form water, so hydrogen is oxidised. This reaction is redox because oxygen is gained by one reactant and lost by another. Do not describe the entire reaction as only oxidation or only reduction. Instead, identify which substance undergoes each change. Following the oxygen makes those roles clear without needing to memorise the example. The name reduction does not mean that the substance simply becomes smaller or lighter.

2:25It describes a particular chemical change, here the loss of oxygen. Likewise, oxidation is not just another word for reacting with air. The oxygen in this example comes from copper(II) oxide, not from the surrounding atmosphere. A second Core example is magnesium burning in oxygen to form magnesium oxide. Magnesium gains oxygen, so magnesium is oxidised. With only the oxygen definitions, no substance here visibly loses oxygen. Later, electrons and oxidation numbers will show that the oxygen is reduced at the same time, so this is still redox.

2:57Some redox reactions contain no oxygen at all. The oxygen definitions remain useful, but they cannot identify every example.

Electron transfer in redox

3:04This next part is for the Extended paper. We will use electrons and oxidation numbers to recognise the same paired changes in a wider range of reactions. Oxidation is loss of electrons, and reduction is gain of electrons. A helpful memory aid is oil rig, which stands for oxidation is loss, reduction is gain. A redox reaction therefore involves both electron loss and electron gain. The electrons lost by one species are gained by another, so oxidation and reduction occur together.

3:34Consider zinc reacting with copper ions that have a 2+ charge. Zinc atoms lose electrons and become zinc ions with a 2+ charge. The copper ions gain those electrons and become neutral copper atoms. This is redox even though oxygen does not appear anywhere in the reaction. The first half equation shows one zinc atom losing two electrons. Losing negative charge leaves the zinc ion with a 2+ charge. In the second half equation, one copper ion gains two electrons.

4:00The two negative electrons cancel its 2+ charge, producing a neutral copper atom. A half equation follows one side of the electron transfer. When we combine the two halves, the electrons cancel because they are transferred rather than created or destroyed. Two leave zinc, and two enter the copper ion. The overall equation therefore shows both chemical changes without displaying the transferred electrons. A common mistake is to assume that gaining electrons makes a particle more positive.

4:30Electrons have negative charge, so gaining them makes a particle more negative, or less positive. Losing electrons has the opposite effect. Thinking about their charge is more reliable than guessing from the words gain and loss. Oxidation numbers provide another way to follow these changes. Oxidation is an increase in oxidation number, and reduction is a decrease. These definitions let us identify redox from an equation even when it does not explicitly show electrons.

4:56In this reaction, zinc changes from zero to positive two. That is an increase, matching the loss of two electrons. Copper goes from positive two to zero, which is a decrease. Its reduction agrees with the gain of two electrons, so both descriptions identify the same changes.

Calculating oxidation numbers

5:12To assign oxidation numbers correctly, start with four rules. The first says that an element in its uncombined state has oxidation number zero. Uncombined means it is present as the element, rather than chemically combined with another element. Zinc metal, copper metal and hydrogen gas therefore all have oxidation number zero. An element can still have oxidation number zero when its particles contain several atoms. Hydrogen gas contains molecules made from two hydrogen atoms, but it remains an uncombined element.

5:40Each hydrogen atom therefore has oxidation number zero. Being joined to another atom of the same element does not make it a compound. The second rule concerns a monatomic ion, which is an ion containing only one atom. Its oxidation number is the same as its charge. A zinc ion with a 2+ charge has oxidation number positive two. For a chloride ion with a 1− charge, the oxidation number is negative one.

6:05The third rule says that the sum of the oxidation numbers in a compound is zero. Here, compound means an electrically neutral substance. Count every atom when making the total, because each contributes its oxidation number. A formula with three oxygen atoms includes three oxygen contributions, not just one. For the examples here, oxygen has oxidation number negative two in compounds. Hydrogen has positive one when combined with non-metals, as in water.

6:31These usual values give us a starting point for finding an unknown oxidation number. They are not universal rules for every possible substance, and an uncombined element still has value zero. Let us find the oxidation number of iron in the oxide with the formula on the board. There are three oxygen atoms, each assigned negative two, giving a total of negative six. The compound is neutral, so the iron contribution must total positive six.

6:54This balances the oxygen contribution and gives an overall sum of zero. There are two iron atoms sharing that total contribution of positive six. Dividing positive six by two gives positive three for each iron atom. The compound is therefore named iron(III) oxide, with three written as a Roman numeral. Notice how the formula determines the value rather than the numeral counting atoms. The fourth rule applies to an ion containing more than one atom.

7:20The sum of its oxidation numbers equals the charge on the ion. Unlike a neutral compound, this total is not necessarily zero. For a sulfate ion with a 2− charge, all the contributions must add to negative two. In the sulfate ion, four oxygen atoms contribute four times negative two, giving negative eight. The whole ion has a 2− charge, so sulfur plus negative eight must equal negative two. Adding eight gives sulfur an oxidation number of positive six.

7:47Using zero as the total would give the wrong answer because this particle is charged. These two calculations use the same balancing idea, but their target totals differ. Iron(III) oxide is neutral, so its total is zero. The sulfate ion carries a 2− charge, making its total negative two. Always check whether the formula represents a neutral compound or an ion before calculating.

Identifying redox and agents

8:10We can now check the original oxygen-transfer reaction using oxidation numbers. Copper changes from positive two in copper(II) oxide to zero in the metal. Its oxidation number decreases, confirming reduction. Hydrogen changes from zero in the element to positive one in water, confirming oxidation through an increase. Oxygen remains at negative two on both sides of this equation. Transferring oxygen does not mean that oxygen itself must change oxidation number.

8:35Here, copper and hydrogen undergo those changes. This example connects the oxygen definitions with the oxidation-number definitions without treating every atom as oxidised or reduced. To identify redox from oxidation numbers, compare the same element before and after the reaction. An increase identifies oxidation, while a decrease identifies reduction. Changes in formula alone are not enough, because atoms can rearrange without changing oxidation numbers. Look for the paired changes rather than assuming that every chemical reaction is redox.

9:04Neutralisation gives a useful contrast. Sodium hydroxide plus hydrochloric acid gives sodium chloride plus water. Sodium stays at positive one, chlorine remains at negative one, and oxygen keeps negative two throughout. Hydrogen is positive one on both sides, so no element changes its oxidation number and this reaction is not redox. An oxidising agent is a substance that oxidises another substance and is itself reduced. In contrast, a reducing agent is a substance that reduces another substance and is itself oxidised.

9:37These names describe what the agent causes in its reaction partner. That is why its own change sounds opposite to its name. In our example, copper ions accept electrons from zinc. By taking those electrons, they cause zinc to lose them and become oxidised. The copper ions are therefore the oxidising agent, and they are themselves reduced. Identify the copper ions, not the copper metal produced afterwards, because the agent is the reacting species.

10:02Zinc supplies the electrons that reduce the copper ions. It is therefore the reducing agent, and zinc itself is oxidised. A useful check is to follow the electrons rather than the name alone. The electron supplier is the reducing agent, while the electron acceptor is the oxidising agent. The same reasoning also applies to the oxygen-transfer example. Hydrogen removes oxygen from copper(II) oxide, so hydrogen acts as the reducing agent.

10:27Copper(II) oxide supplies oxygen to hydrogen and is itself reduced. It therefore acts as the oxidising agent, consistent with the definitions we have just established.

Redox tests and exam answers

10:38Some redox reactions produce useful colour changes. Acidified aqueous potassium manganate(VII) is purple. When it is reduced, the solution changes from purple to colourless. This reagent therefore detects a reducing agent, because that substance causes the manganate(VII) to undergo reduction. The Roman numeral seven identifies the oxidation number of manganese in the manganate ion. During reduction, manganese changes to a lower oxidation number. That chemical change produces the observed loss of purple colour.

11:05Include the word acidified when you name this reagent, because the test uses it in acidic solution. Separate the observation from its interpretation. Purple to colourless describes what you see. Stating that the substance tested is a reducing agent explains what the result means. Do not call it an oxidising agent simply because the reagent has undergone a colour change. Aqueous potassium iodide provides a complementary test. Its solution changes from colourless to yellow-brown when iodide ions are oxidised to iodine.

11:34An oxidising agent causes that change by accepting electrons from iodide ions. This time, the substance being tested is identified as an oxidising agent rather than a reducing agent. Each iodide ion loses one electron during this process. Two iodide ions together form one iodine molecule and release two electrons. Iodine changes from oxidation number negative one to zero, which is an increase. That confirms oxidation, even though the final value is zero rather than a positive number.

12:03These observations can appear in practical or alternative to practical questions. Record both the starting colour and the final colour when asked to describe the change. A separate question may ask you to explain the result using oxidation or reduction. A colour change alone does not prove redox in general, but these two tests are standard redox tests. An exam question might show the zinc equation and ask you to state what is oxidised.

12:25It could then ask you to explain why the reaction is redox and identify the oxidising agent. Begin by assigning oxidation numbers or tracking electrons. Both methods lead to the same answer, but your explanation must name the relevant species. Zinc is oxidised because it loses electrons. The copper ions gain those electrons, so they are reduced. The reaction is redox because oxidation and reduction occur simultaneously. Copper ions are the oxidising agent because they oxidise zinc and are themselves reduced.

12:55If the question instead asks for an explanation using oxidation numbers, give those values explicitly. The oxidation number of zinc increases from zero to positive two, so zinc is oxidised. Copper is reduced, because its value decreases from positive two to zero. These paired changes show that the reaction is redox. For the colour-test question shown, first state that the substance is a reducing agent. Acidified potassium manganate(VII) is reduced, producing the change from purple to colourless.

13:24The substance tested causes that reduction and is itself oxidised. If asked only to describe the observation, give the colour change rather than replacing it with an explanation. Three mistakes are especially important to avoid. A Roman numeral gives an oxidation number, not an atom count. Reduction is gain of electrons, which makes a particle less positive, not more positive. An oxidising agent is itself reduced, since its name describes the change it causes in another substance.

13:51Redox is one connected idea viewed in three ways. Oxygen transfer involves gain and loss of oxygen, while electron transfer involves simultaneous loss and gain of electrons. Oxidation numbers reveal those changes through an increase and a decrease. Whichever method you use, follow both partners and keep their roles separate. Use the four oxidation-number rules to establish values before comparing reactants with products. Identify the oxidising agent as the substance that is reduced, and the reducing agent as the substance that is oxidised.

14:20For the tests, remember purple to colourless with acidified potassium manganate(VII), and colourless to yellow-brown with potassium iodide. Each observation now connects to a chemical change you can explain.