Have you ever wondered how batteries or cell phone batteries work? Or why cell phone batteries can be recharged while batteries can only be used once? This is due to redox reactions.
The redox reaction and its components
A redox reaction is an electron transfer reaction. The name comes from the two partial reactions into which the redox reaction can be divided: Oxidentification and editorauction. Both partial reactions take place simultaneously.
redox reaction – oxidation
An oxidation is defined as a reaction in which electrons are lost from a substance. It is said that the substance, which also acts as a reducing agent, is oxidized. Oxidation increases the oxidation number.
In general: A substance A releases electrons during oxidation.
Originally, the term «oxidation» was influenced by Antoine Laurent de Lavoisier. He used the term for oxide formation, i.e. the combination of an element or a molecule with elementary oxygen. Later, in addition to oxide formation, dehydrogenation was also referred to as oxidation. Dehydrogenation is the removal of hydrogen atoms from a molecule. In the meantime, the term has been generalized by further chemical knowledge (such as Bohr’s atomic model and the ion theory).
redox reaction – reduction
A reduction is defined as a reaction in which electrons are accepted from a substance. It is said that the substance, which is also called the oxidizing agent, is reduced. The reduction decreases the oxidation number.
In general: A substance B takes up electrons during reduction.
Reduction used to be considered the reverse of oxidation, so at that time a reverse of oxide formation. Reduction comes from the Latin word reductio and means something like «return».
oxidizing agent
In a redox reaction, the oxidizing agent is reduced while oxidizing another substance. Oxidizing agents accept electrons during reduction, which is why they are also called electron acceptors.
The elements of the main group of chalcogens (main group 6) and halogens (main group 7) are designated as strong oxidizing agents. These include, for example, the chalcogen oxygen (O) and the halogens fluorine (F) and chlorine (Cl).
reducing agent
The reducing agent is oxidized in the redox reaction while the oxidizing agent is reduced. Reducing agents give off electrons during oxidation, which is why they are also called electron donors.
These include above all alkali metals and alkaline earth metals, i.e. chemical elements of the 1st and 2nd main groups. In addition to sodium (1st main group) and magnesium (2nd main group), other base metals such as aluminum and zinc are also good reducing agents.
redox reaction
In summary, the oxidation and the reduction result in the following overall equation:
Since there is an electron on both sides, this can be shortened from the equation and thus we get:
The donor-acceptor principle generally describes the transfer of particles during a chemical reaction. Particles are transferred from one reactant to another.
The reactant that gives off particles is called the donor, while the reactant that takes in particles is called the acceptor. An example of reactions that work according to the donor-acceptor principle are acid-base reactions. In this process, protons are transferred between the proton donor and the proton acceptor.
Redox reactions also run according to the donor-acceptor principle. The oxidizing agent gains electrons while the reducing agent gives up the electrons.
redox reaction – oxidation numbers
Using the oxidation numbers, you can divide a redox equation into its sub-equations (reduction and oxidation).
The oxidation numbers indicate the (partial) charge of an atom in a molecule/ion. The electronegativity of the individual atoms is compared. Oxidation increases the oxidation number by losing electrons, while reduction decreases the oxidation number by gaining electrons.
- The sum of the oxidation numbers of the atoms in a molecule must result in the molecule or ion charge.
- Chemical elements or element molecules (such as elemental hydrogen (H2)) always have the oxidation number 0.
- If it is a monatomic ion, the oxidation number is equal to the charge of the ion (examples: Cl-, O2-, …).
- In covalent compounds, one compares the electronegativity of the bonding partners and assigns the bonding electrons to the atom that is more electronegative.
If you want to learn more about oxidation numbers, you can read the article on this topic.
redox reaction – set up reaction equations
You can always use the same principle to set up a reaction equation for a redox reaction. The following steps are necessary for this:
- Determination of all oxidation numbers
- Classification of partial reactions in oxidation and reduction
- Add donated and accepted electrons to the partial reactions
- Charge balancing with the help of hydroxide ions (OH-) in an alkaline environment or oxonium ions (H3O+) in a neutral or acidic environment
- Mass balance with the help of water molecules (due to step 4)
- balance of electrons
- Merging the partial equations and shortening the overall equation
It is important to know that not all steps are mandatory. In some redox reactions, steps can be skipped. Among other things, one does not need to balance the charge if the overall charge of the reaction is already neutral. Therefore step 5 is also not necessary. Step 6 is not necessary if both partial reactions already have the same number of electrons. The other steps are mandatory.
Setting up a redox reaction – example
This reaction equation describes a redox reaction that takes place during the ignition of fireworks: Potassium chlorate (KClO3) reacts with sulfur (S) to form potassium chloride (KCl) and sulfur dioxide (SO2).
Proceed as explained above and start by determining the oxidation numbers.
1st step: Determination of all oxidation numbers
You start by determining the simplest oxidation numbers:
- Sulfur occurs as an element on the educt side (left-hand side) of the reaction equation, which is why it has an oxidation number of 0.
- Since oxygen does not occur in combination with fluorine or as a peroxide, every oxygen atom in the entire reaction equation has the oxidation number -II.
- Potassium is in the 1st main group in the periodic table, so it belongs to the alkali metals. Therefore, throughout the reaction equation, potassium has the oxidation number +I.
- Since potassium chlorate (KClO3), as a neutral molecule, must have a sum of all oxidation numbers of 0, the oxidation number of chlorine can be calculated. A potassium chlorate molecule consists of one potassium atom, one chlorine atom and three oxygen atoms. From this one can form the following equation:
- This gives an oxidation number of +V for chlorine.
- The same applies to potassium chloride (KCl):
- In contrast to potassium chlorate, chlorine has the oxidation number -1.
- Now the oxidation number of sulfur in sulfur dioxide (SO2) has to be calculated:
- The oxidation number for sulfur in sulfur dioxide is therefore +IV.
2nd step: Classification of partial reactions in oxidation and reduction
While the oxidation numbers of oxygen and potassium do not change, the oxidation number of sulfur increases. On the left side of the reaction equation it has an oxidation number of 0 and on the right side an oxidation number of +IV. That is, sulfur is oxidized. On the other hand, chlorine has an oxidation number of +V on the left-hand side of the reaction equation, while it only has an oxidation number of -I on the right-hand side. So chlorine is reduced.
3rd step: add the given and accepted electrons to the partial reactions
Now add the difference in the oxidation numbers as electrons gained or lost to the partial equations:
4th step: charge balancing
Oxonium ions (H3O+) are used here for charge equalization. Since there are four negatively charged electrons during oxidation, we add four positively charged oxonium ions to balance the charge. In the reduction, there are six electrons on the left side of the reaction equation, so we add six more oxonium ions there.
5th step: material equalization
The matter of balancing is to add water molecules to the partial equations so that the number of water and oxygen atoms is balanced again.
In the oxidation, there are twelve hydrogen atoms and six oxygen atoms on the right. There are no hydrogen atoms on the left and no oxygen atoms either. So you add six water molecules (H20) to this side, so that you get a total of eighteen hydrogen atoms and nine oxygen atoms on this side.
You can use the same principle for the reduction. Nine water molecules (H2O) are added to the right-hand side, giving eighteen hydrogen atoms and new oxygen atoms on both sides of the reaction equation.
6th step: Balancing the electrons
Now you have to complete both partial equations in such a way that both have the same number of electrons.
7th step: merging the partial equations and shortening the overall equation
Now you can put the two sub-equations together.
Then you cut out the molecules from the equation that appear in the same number on both sides.
In summary, potassium chlorate (KClO3) is reduced to potassium chloride (KCl) while sulfur (S) is oxidized to sulfur oxide (SO2). That is, while potassium chlorate gains electrons, sulfur loses electrons.
Redox reaction – examples
electrolysis
In chemistry, electrolysis is the process in which a redox reaction is forced by an electric current. For example, electrolysis is used to split water into hydrogen and oxygen. In addition, electrolysis is used to extract metals such as aluminum.
Basically, electrolysis is the reverse of what happens in a battery. Instead of converting chemical energy into electrical energy, exactly the opposite happens in electrolysis: electrical current is converted into chemical energy.
During electrolysis, a voltage is applied to the two electrodes of the electrolytic cell. The electrodes, anode (= positive pole) and cathode (= negative pole), are immersed in an ionic solution, also known as an electrolyte solution. The applied voltage causes electrons to flow to the cathode, which is connected to the positive pole. An excess of electrons is formed, which is why the cathode acts as a negative pole during electrolysis.
At the anode, in turn, the electrode that is connected to the negative pole, there is a lack of electrons. Due to the excess of electrons at the cathode, positive ions, so-called cations, are attracted. They absorb the excess electrons at the cathode. So they are reduced. The negatively charged anions are attracted to the anode and give up their electrons there. They are oxidized. The number of…