Have you ever wondered why an apple slowly turns brown when it is cut into two and exposed to air? Do you know that lemon juice is able to prevent such a process? These occurrences involve the natural phenomena called reduction and oxidation that occur simultaneously; this is better known as a redox reaction. There are many redox reactions occur in our life. The process of inhaling oxygen and exhaling carbon dioxide during the breathing process is also an example of a redox reaction. This section will introduce to you the concept of redox reactions.
- The redox reactionIf we divide the word "redox" into two parts "red" and "ox", the "red" comes from the word "reduction" while the "ox" comes from the word "oxidation". Thus, the word "redox" actually means the reduction-oxidation process. In other words, a redox reaction involves both the oxidation and reduction processes that occur simultaneously. There are many processes in everyday life involving redox reactions. These include rusting, decomposition, respiration, etc.We shall look at the oxidation reaction followed by the reduction reaction. These two reactions can be explained based on the four following processes:i. Loss or gain of electronsii. Loss or gain of hydrogeniii. Transfer of electronsiv. Change in the oxidation numberOxidation can be defined as the gain of oxygen, the loss of hydrogen, the loss of electrons or an increase in the oxidation number of the element in the compound. We shall each one of these in detail.
- Gain of oxygenWhen a substance gains oxygen, it is said to be oxidised. For example,In this case, calcium burns in the excess of oxygen to form calcium oxide. Calcium is said to be oxidised to form calcium oxide as it gains oxygen in the reaction.Loss of hydrogenNot all oxidation reactions involve oxygen. If a substance loses hydrogen during a reaction, it is also said to be oxidised. Let us refer to the reaction below,In the above reaction, ammonia loses its hydrogen to form nitrogen gas. Thus, ammonia is said to be oxidised to nitrogen gas.Loss of electronsFor the reactions that do not involve oxygen and hydrogen, we can explain the reaction based on the transfer of electrons. For instance, the formation of sodium chloride involves the transfer of electrons from sodium to chlorine,ReductionReduction is the reverse of oxidation. It can be defined as the loss of oxygen, the gain of hydrogen, the gain of electrons or the decrease in the oxidation number of the element.Loss of oxygenMagnesium oxide reacts with carbon to eliminate oxygen. Magnesium oxide is said to be reduced to magnesium because it loses oxygen in the reaction.Gain of hydrogenAs in oxidation, not all reduction reactions involve oxygen. If a substance gains hydrogen during a reaction, it is also said to be reduced.Gain of electronsAs in oxidation, reduction can also be defined based on the transfer of electrons.The chlorine gas accepts electrons to form chlorine ions. Thus, chlorine gas is said to be reduced to chloride ion.Decrease in the oxidation numberThe reduction can also be explained using the oxidation number. The reduction decreases in the oxidation number.Redox Reactions Based On Changes In The Oxidation NumberFrom the discussion above, we can see that there are four techniques to define the process of oxidation and reduction. There are the gain (or loss) of oxygen, the gain (or loss) of hydrogen, the transfer of electrons and the change in the oxidation state of the element. Among the four techniques, the change of the oxidation state is comprehensively used. Accordingly, we need to discuss this technique in detail. Let us begin with an understanding of the concept of the oxidation number.The oxidation numberThe oxidation number is also known as the oxidation state. It is the imaginary charge of an element as if it exists as an ion. The oxidation number of the element can be assigned according to a set of rules.Rules for assigning the oxidation numberRule 1An atom or a molecule in its elemental state has an oxidation number of zero. For example, Na, Mg, Al, C, O2, H2, Cl2, etc. each has the oxidation number of zero.Rule 2The oxidation number of a mono-atomic ion is equal to the charge carried by the ion.Rule 3The element which is more electronegative is given a negative oxidation number.a) The oxidation number of fluorine is always -1 because it is the most electronegative element in nature.b) For other halogens, the oxidation number in a compound is always -1 as well unless it combines with more electronegative elements.c) The oxidation number of oxygen in a compound is always -2 except in peroxide and when it combines with other more electronegative elements.d) The oxidation state of hydrogen in a compound is always +1unless it combines with reactive metals to form metal hydrides,Rule 4The sum of the oxidation numbers of all the elements in the formula of a compound is equal to zero. This rule can be applied for both ionic and covalent compounds.Rule 5The sum of the oxidation number of all the elements in the formula of the polyatomic ion must be equal to the charge of the ion.For example,
- Half- EquationsIn representing redox reactions, we need to know the concept of half-equations or half-reactions. A complete redox reaction involves two reactions, the oxidising reaction and the reducing reaction. We can therefore say that oxidising is carried out by half of the reaction while the reducing reaction is carried our by the other half. The equation representing only the oxidising or only the reducing process is called a half- reaction or a half-equation.
The Oxidation Number And Nomenclature Of CompoundsOne of the characteristics for transition elements is that they can exist in more than one oxidation state. Thus, to avoid confusion, the International Union of Pure and Applied Chemistry (IUPAC) system is used to include the oxidation number in the naming of inorganic compounds.By referring to the names of some simple ionic compounds which have transition metals as cations, the oxidation number is included in the IUPAC nomenclature. The oxidation number of a metal ion is represented by a Roman numeral in brackets, immediately followed by the name of the metal. The table below illustrates some examples:The same principle is applied to the naming of anions. The oxidation number of a metal in the anion is represented by a Roman numeral in brackets after the name of the metal. One important point we need to take note of here is that for anions, the name of the metal ends with "ate". The table below illustrates some examples.Besides transition metals, some non-elements also exhibit more than one oxidation number. The same principle is applied. Let us look at the examples below:Lastly, we have to bear in mind that there is no such name as sodium (I) chloride, magnesium (II) oxide, etc. We have to understand that elements from Groups 1, 2 and 13 in the periodic table only exhibit one oxidation number. Thus, there is no Roman numeral written in brackets after the name of the metals.Redox Reactions Based On Changes In The Oxidation StateWe know now that when oxidation occurs, there is an increase in the oxidation number; conversely, when reduction occurs, there is a decrease in the oxidation number. The following example explains the concept,Since the oxidation number of bromine in hydrobromic acid increases from -1 to zero, the hydrobromic acid is said to be oxidised to bromine. At the same time, the chlorine is said to be reduced to hydrochloric acid as its oxidation number has decreased from zero to -1.Take note that not all the reactions give the change of the oxidation number of elements in the reactants. For example, consider the reaction of neutralisation,There is no change for all the elements during the reaction. Hence, it is not a redox reaction.Oxidising And Reducing AgentsIn a redox reaction, oxidation and reduction occur simultaneously; therefore, a substance will be oxidised and another substance will be reduced at the same time. Here, we introduce two new terminologies: the oxidising agent and the reducing agent. An oxidising agent is a substance that oxidises or causes the oxidation of another substance in a redox reaction. It is itself reduced in the reaction. A reducing agent is a substance that reduces another substance in a redox reaction. It is itself oxidised in the reaction.Let us consider the reaction between copper (II) oxide and carbon,Copper (II) oxide oxidises carbon to form carbon dioxide while it itself experiences reduction. Thus, it is an oxidising agent. If we look from the perspective of carbon, we find that it reduces copper (II) oxide to copper. Hence, it is a reducing agent.The table below shows some examples of oxidising and reducing agents.Note the following characteristics of reducing and oxidising agents:i. One of the reactants undergoes oxidation (indicated by the increase in the oxidation state of an element) - the reducing agent.ii. The other reactant undergoes reduction (indicated by the decrease in the oxidation state of an element) - the oxidising agent.Types Of Redox ReactionsAs mentioned above, not all reactions are redox reactions. Here, we are going to show you some examples of redox reactions and then we take note of a few characteristics of redox reactions. The following are some examples of redox reactions.Combustion of a metal in oxygenA metal combines with oxygen to form metal oxide.
- Displacement of metals from their salt solutionMetals are electropositive. They tend to donate a electron or electrons to form positive ions. In other words, they tend to undergo oxidation and therefore act as reducing agents. Every metal has a different electropositivity and thus, a different reducing ability. The electropositivity of metals is listed in the electrochemical series (ECS) as follows:
The electrochemical seriesThe higher the position of a metal in the electrochemical series, the more electropositive the metal is and the easier it is for it to donate electrons. Hence, the better reducing agent it is. Thus, when a metal at a higher position in the ECS is mixed with metal ions at a lower position in the ECS, it has a higher tendency to form positive ions and therefore reduces the other metal ions to precipitate. This phenomenon can be concluded by one rule,Let us consider the example below.Displacement of halogens from their halide solutionsHalogens are electronegative. They tend to receive a electron or electrons to form negative ions. In other words, they tend to undergo reduction and therefore act as oxidising agents. When going down the group, the number of occupied shells increases. Since the outermost occupied shell is further away from the nucleus, the ability to attract electrons to form negative ions decreases down the group. Hence, the electronegativity decreases and the ability for a halogen to be an oxidising agent decreases.
The oxidising ability of halogensSimilar to the displacement of metals, there is a rule for the displacement of halogens,The principle behind this is that a more electronegative halogen has a higher tendency to form negative ions. This natural tendency leads to a reaction when it is mixed with the ion of a halogen which is less electronegative. Let us consider the example below:The presence of the halogen can be confirmed by using an organic solvent like 1,1,1-trichloroethane. The table below shows the colour exhibited by halogens in a aqueous solution and 1,1,1-trichlororethane.Heating metallic oxides with carbonIn the industry, carbon, in the form of coke, is used to extract metals from its ore in conditions when the metal is located below it in the reactivity series. For example, let us consider the extraction of iron from iron ore.We will study about reactivity series in more detail later.Transfer of electrons at a distanceA redox reaction can still occur even if the oxidising agent is separated from the reducing agent. This is carried out by the transfer of the electrons at a distance. See the experimental set-up below:
A redox reaction by the transfer of electrons at a distanceThe acidified potassium manganate (VII) KMNO4 is put at one arm of the U-tube and the iron (II) sulphate FeSO4 at the other arm. The Fe2+ ions of the FeSO4 solution act as the reducing agent, releasing electrons to form Fe3+ ions. Hence, the colour of the solution changes from pale green to yellow. The oxidation half- equation is,The electrons then flow through the external wire to reach another arm. The needle of the galvanometer is deflected. The electrons reaching the arm are received by the manganate (VII) ions MnO4- of KMnO4 solution. The purple colour ions are reduced to manganese (VII) ions Mn2+ which is colourless. The reduction half- equation is,The overall ionic equation is as follows,Take note that this is a reaction for the electrochemical cell. The electrode at which the electrons are released by the reducing agent is called the negative terminal while another electrode at which the electrons are received is called the positive terminal.The reactivity seriesThe reactivity series is a list of elements arranged in terms of their reactivity. The higher the element in the list, the more reactive is the element and the easier it is for this element to lose electrons and form positive ions.