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Redox reactions or oxide reduction reactions are chemical processes in which net transfers of electrons occur from one chemical species that is oxidized to another that is reduced . This type of reactions is difficult to adjust by traditional methods such as trial and error, so alternative methods have been developed that facilitate the process. One of these methods is the half-reaction method, also known as the electron ion method .
What is the method of half-reactions or the electron ion?
The half-reaction method consists of a set of steps to follow to balance or adjust the equations of redox reactions. This method is based on the idea that redox processes actually consist of the coupling of two processes that can be considered separately, which are oxidation and reduction.
In the method of half-reactions or the method of the electron ion, the equations of the oxidation and reduction half-reactions are adjusted separately to later combine both equations in an already balanced global equation.
The oxidation and reduction half-reactions
Oxidation is a chemical process during which an atom or a group of atoms loses or releases one or more electrons . This process necessarily implies an increase in the oxidation state of some of the atoms that make up the original species.
On the other hand, reduction is understood as the opposite process to oxidation. Reduction is the chemical process during which a chemical species gains one or more electrons . When this happens, the oxidation state of some of the atoms that make up this chemical species decreases, since it is receiving an electron whose charge is negative.
Two halves of the same process
Free electrons are extremely unstable species, so the oxidation reaction is a process that cannot occur independently, except under very particular conditions. In other words, it cannot happen that an atom spontaneously releases an electron without further ado, and that this electron remains, so to speak, “floating around”. This only occurs in highly energetic conditions, such as in plasma, or when a material is bombarded with some type of high energy radiation. Consequently, oxidation reactions can only occur if at the same time another species is capable of receiving the released electrons.
In view of this, oxidation and reduction cannot be considered as chemical reactions in themselves, but rather are two halves of the same process, which is why they are called half-reactions or half-reactions, although the latter The term is rarely used in the Spanish chemical literature.
The half-reaction method to adjust redox reactions
Next, the steps to balance the equation of a redox reaction using the electron ion method or half-reaction method will be detailed.
It should be noted that this method admits two variants depending on whether the reaction is carried out in an acid medium or in a basic medium. In much of the literature these two methods are detailed separately, following slightly different steps during different stages of the process. However, a redox-adjusted reaction in an acidic medium can easily be converted to a basic medium by means of three very simple steps. For this reason, we think it is more convenient to learn how to set up reactions in an acid medium (which is easier) and then transform it to a basic medium if necessary.
To illustrate this process, we will fit the following redox reaction that occurs in a basic medium:
Step 0 (optional): Dissociate all dissolved ionic species to obtain the ionic equation
The adjustment process by the electron ion method is much simpler if all spectator ions are excluded from the half-reactions, that is, all those ions that are not directly involved in the oxidation or reduction but are nonetheless present in the reaction. solution and form part of the original ionic compounds.
The first step in doing so is to dissociate all dissolved ionic species, that is, salts, acids, and bases. Those ions that appear on both sides of the equation completely unchanged will be the spectator ions. In the case of our example, the ionic equation will be like this:
Looking at this equation, it is clear that the potassium cation is not involved in the reaction and is therefore a spectator ion. Then, the net ionic equation that we will adjust, after eliminating this ion, will be:
This step is not always necessary, since in some cases we start directly from the net ionic equation (the one in which the spectator ions are no longer present), and in others, the equation is so simple that the presence of these ions does not interfere in the reaction adjustment process.
Step 1: Identify the species that are being oxidized and reduced.
The next step involves determining the oxidation state of all the atoms present in the chemical equation, in order to know which atoms underwent a change in oxidation state. There must necessarily be at least one atom that is oxidized and one that is reduced, and it may even be the same atom (in which case we are in the presence of a particular type of redox reaction called dismutation).
It is not the purpose of this article to give a complete explanation on how to determine the oxidation states, but let’s remember as basic rules that:
- Elemental substances have oxidation state 0.
- The oxidation state of monatomic cations and anions corresponds to their charge.
- In all oxides and oxyanions, oxygen has -2 oxidation states.
- With the exception of hydrides, where its oxidation state is -1, hydrogen always has a +1 oxidation state in all the compounds of which it is a part.
- The other oxidation states are calculated in such a way that the sum of all oxidation states matches the net charge of the species in question.
The following equation presents the oxidation states of all the species involved in our example:
As we can see, the atoms that are changing oxidation states are manganese and iodine. The manganese in the permanganate ion is being reduced from +7 to +4 while the iodide is being oxidized to elemental iodine, going from -1 to 0 oxidation state.
Step 2: Separate the overall reaction into oxidation and reduction half-reactions.
Now that we know which species are being oxidized and reduced, we can divide the overall reaction into two half-reactions:
Note that, since hydroxide ions are not directly involved in the oxidation or reduction process, they were not included in any of the half-reactions.
Step 3: Separately equilibrate the two half-reactions as if they were in an acid medium.
As explained at the beginning, whether the reaction occurs in an acid medium or if it is basic, we will begin adjusting it as if it occurred in an acid medium. Later, if necessary, it will be transformed into a basic medium. The adjustment of the half-reactions in acid medium consists of the following 5 steps, which can be applied simultaneously to both half-reactions:
- Adjust the number of atoms that are changing oxidation states.
In our case, the reduction does not cause any change, since there is one manganese on each side, but the oxidation does:
- Adjust for anything other than oxygen or hydrogen, adding spectator ions if necessary.
In our example this is not necessary, since we remove all spectator ions at the beginning.
- Adjust the number of oxygens by adding water molecules where they are missing.
In our case, it is necessary to adjust the number of oxygens in the reduction half-reaction, but not in the oxidation:
- Adjust the number of hydrogens by adding protons (H + ) where they are missing:
Again, the oxidation remains unchanged because it does not involve hydrogen atoms, but in the reduction we do need to adjust them:
- Adjust the total electric charge by adding electrons (e – ) where there are missing negative charges or excess positive charges (Tip: they are almost always on the same side as the protons):
As can be seen, in the reduction half-reaction the net charge on the products is 0, but on the reactants there is a net charge of +4 – 1 = +3, that is, there are surplus positive charges. For this reason, we must add three electrons on the side of the reactants to compensate for this excess charge:
On the other hand, in the case of oxidation, there is a net charge of –2 on the reactants side and 0 on the products, so there are no negative charges on the products, so 2 electrons must be added on this side to balance the charges:
Clue
It should be noted that the addition of electrons by this procedure (treating them as if they were ions, hence the name of the ion-electron method) is done independently of the oxidation states of the different species involved. However, it is essential that the number of electrons and their placement match the observed changes in oxidation states.
Thus, in reduction half-reactions, electrons must always be on the left side of the equation and in oxidations they must always be on the right side, as happened in our example.
Also, the number of electrons must match the change in oxidation state. Manganese is reduced from +7 to +4, so there is a -3 change in its oxidation state, consistent with the addition of 3 electrons. In the case of iodide, this changes from -1 to 0 corresponding to a change of +1, but there are two iodides, so two electrons are released instead of one, as presented in the respective equation.
Step 4: Multiply each half-reaction by the number of electrons in the other, simplifying the factors if possible.
This step seeks to equalize the number of electrons released during oxidation with the number of electrons captured by reduction. This ensures that there are no “orphan” electrons at the end of the reaction or that no electrons are missing. If both half-reactions release or take up the same number of electrons, this step is not necessary.
In our example, each oxidation half-reaction releases 2 electrons, but each reduction half-reaction requires 3, so oxidation needs to occur 3 times for every 2 times reduction occurs:
The result is:
Step 5: Add both half-reactions to get the balanced net ionic equation.
The sum of these two half-reactions results in the adjusted net ionic equation in an acid medium:
Step 6 (for basic medium only): Convert the acidic medium to a basic medium.
At the end of step 5 we already have the adjusted net ionic equation in an acid medium. However, the reaction may occur in a basic rather than an acid medium. If this is the case, the previous equation must be transformed to a basic medium. This is done through three simple steps:
- Add one hydroxide ion (OH – ) to each side of the equation for each proton (H + ) present.
In our case, 8 hydroxide ions must be added from each side:
- Combine the hydroxides and protons that are on the same side to form water molecules.
In our case, in the reactants there are 8 hydroxides and 8 protons that are neutralized to form 8 water molecules:
- If necessary, simplify the water molecules that are repeated on both sides of the equation.
This last step results in the balanced net ionic equation in basic medium. In the case of the reaction that we are adjusting, after forming the 8 water molecules, we can notice that only four of these eight actually participate in the reaction, since the other four remain unchanged in the products. Simplifying these four repeating water molecules gives the adjusted redox equation:
Step 7 (optional): Add the spectator ions to obtain the overall molecular equation
This step is not always necessary, since the net ionic equation is a more accurate representation of the chemical process that is actually occurring. However, it can be important for carrying out stoichiometric calculations. In this sense, if you want to obtain the global molecular equation, you only need to add the spectator ions as counterions of all the species that appear in the net ionic equation.
In the present example, the only spectator ion is the potassium cation (K + ), so we will use it to neutralize all the anions present in the reaction:
Finally, after uniting the respective ions, we obtain the adjusted equation in terms of neutral species only:
References
Chang, R., & Goldsby, K. (2013). Chemistry (11th ed.). McGraw-Hill Interamericana de España SL
Generalic, E. (2021, January 22). Balancing of redox reactions by the ion-electron method . periodni.com. https://www.periodni.com/en/method_of_semi-reactions.php
Lavado S., A., & Yenque D., JA (2005). Unified procedure to balance redox reactions using the Ion-Electron method . Redalyc. https://www.redalyc.org/pdf/816/81680214.pdf
