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An oxidation–reduction process, often known as a redox reaction, is a reaction in which electrons are transferred between chemical species (the atoms, ions, or molecules involved in the reaction).
Redox reactions may be found all around us: fuel combustion, metal corrosion, and even photosynthesis and cellular respiration entail oxidation and reduction.
The potential charge of an atom if all of its links to other atoms were totally ionic is called the oxidation state.
Objectives In Learning
By using the group number, you may predict the oxidation states of common elements.
Points to Remember
A pure element’s oxidation state is always zero.
A pure ion’s oxidation state is the same as its ionic charge.
The oxidation state of hydrogen is +1, while the oxidation state of oxygen is -2.
The total oxidation states of all the atoms in a neutral molecule must equal zero.
Term Definitions
Reduction is the acquisition of electrons that causes the oxidation state to drop.
oxidation is defined as the loss of electrons that results in an increase in the oxidation state.
The degree of oxidation for an atom in a chemical compound is indicated by its oxidation state, which is the hypothetical charge that an atom would have if all links to atoms of various elements were entirely ionic.
Integers, which can be positive, negative, or zero, are commonly used to indicate oxidation states. In rare circumstances, an element’s average oxidation state is a fraction, such as 8/3 in magnetite (Fe3O4).
The highest known oxidation state is +8 for ruthenium, xenon, osmium, iridium, hassium, and various plutonium complexes, whereas the lowest known oxidation state is 4 for several carbon group elements.
Oxidation States: General Guidelines
A free element (uncombined element) has no oxidation state.
The oxidation state of a simple (monoatomic) ion is equal to the ion’s net charge. The oxidation state of Cl–, for example, is -1.
Hydrogen has an oxidation state of +1 in most molecules, while oxygen has an oxidation state of 2. Hydrogen has an oxidation state of 1 in active metal hydrides (such as LiH), an oxidation state of 1 in peroxides (such as H2O2), and an oxidation state of -1/2 in superoxides (such as KO).
In a neutral molecule, the algebraic total of oxidation states for all atoms must be zero. In ions, the algebraic total of the individual atoms’ oxidation states must equal the ion’s charge.
Identifying Oxidation States and Predicting Oxidation States
The oxidation state of most common elements may usually be deduced from their periodic table group number. The following graph summarises this:
According to the table above, boron (a Group III element) has an oxidation state of +3, while nitrogen (a Group V element) has an oxidation state of -3. Keep in mind that oxidation states can alter, therefore this method should only be used as a rough guide; for example, transition metals don’t follow any strict rules and can have a large range of oxidation states.
The charge of a molecule or polyatomic ion is equal to the total of the oxidation states for all atoms in the molecule or ion, as indicated in rule number four above.
This aids in determining the oxidation state of any one element in a given molecule or ion, providing that all of the other elements’ oxidation states are known. For example, the total charge of a sulfite ion (SO32-) is 2-, and each oxygen is assumed to be in its normal oxidation state of -2. Because sulfite contains three oxygen atoms, it provides 3-2=-6 to the overall charge. For the total charge on sulfite to be 2-: 4-6 = -2, sulphur must have an oxidation state of +4.
It’s important not to mix up an atom’s formal charge with its formal oxidation state, as the two might be quite different (and often are different, in polyatomic ions). The charge on the nitrogen atom in the ammonium ion NH4+, for example, is 1+, yet the formal oxidation state is -3, which is the same as nitrogen in ammonia. The formal charge on the N atom changes between ammonium and ammonia, while its oxidation state does not.
Reduction reaction
A reduction reaction occurs when any species gains electrons in a chemical reaction. The positive charge of the participating species diminishes as electrons are gained, whereas the negative charge of the participating species grows.
2Na + Cl2 = 2Na + Cl–, for example.
oxidation state: 01 -1 oxidation state: 0 0 to +1 -1 oxidation state: 0 0
Chlorine is an example of a reduction reaction because its oxidation state changed from 0 to -1, indicating that it gained an electron.
Conclusion
An oxidation–reduction process, often known as a redox reaction, is a reaction in which electrons are transferred between chemical species (the atoms, ions, or molecules involved in the reaction).
Redox reactions may be found all around us: fuel combustion, metal corrosion, and even photosynthesis and cellular respiration entail oxidation and reduction.
The potential charge of an atom if all of its links to other atoms were totally ionic is called the oxidation state.
