Grade 11 → Classification of elements and periodicity in properties → Periodic trends in properties ↓
Oxidation states
The concept of oxidation state is an essential part of understanding how elements react with each other. It is particularly important in the context of redox reactions, where there is a transfer of electrons. The oxidation state, also known as the oxidation number, helps us keep track of electrons during chemical reactions. This concept is important in explaining various chemical properties and predicting the outcomes of reactions.
What is the oxidation state?
The oxidation state is a number assigned to an element in a chemical compound indicating the number of electrons lost (or that can be considered lost) and gained (or that can be considered gained) by an atom of that element in the compound. The concept is used primarily to determine how electrons are transferred in chemical reactions, especially redox reactions. Generally, the oxidation state is positive for metals (indicating a loss of electrons) and negative for nonmetals (indicating a gain of electrons).
Assigning oxidation state
Determining the oxidation state involves some fundamental rules:
- The oxidation state of an element in its pure form is always zero. For example,
O_2
andN_2
both have oxidation state of 0. - For monatomic ions, the oxidation state is the same as the charge of the ion. For example, the oxidation state of Na+ is +1, and that of Cl- is -1.
- In compounds, the oxidation state of hydrogen is usually +1, and the oxidation state of oxygen is usually -2. Exceptions exist, such as in
H_2O_2
where the oxidation state of oxygen is -1. - In a neutral compound, the sum of the oxidation states must be zero. For example, in H2O, the sum of the oxidation states is (+1) * 2 + (-2) = 0.
- For polyatomic ions, the sum of the oxidation states must equal the charge of the ion. For example, in SO42-, assuming that sulfur has an oxidation state of +6 and oxygen has an oxidation state of -2, the sum is 6 + 4*(-2) = -2.
Understanding through visual examples
Let us use some examples to better understand how oxidation states are determined.
Example 1: Water (H2O
)
H - Oxidation state +1 2 x H = 2(+1) = +2 O - Oxidation state -2 Sum = +2 + (-2) = 0
Water is a neutral molecule, so the sum of the oxidation states is zero.
Example 2: Ammonium ion (NH4+
)
N - Oxidation state -3 H - Oxidation state +1 4 x H = 4(+1) = +4 Total = -3 + 4 = +1
The total oxidation number is 1, which is equal to the charge on the ammonium ion.
Oxidation states of elements depending on location in the periodic table
Oxidation states are affected by the element's location in the periodic table, its group, and its period. Different groups of elements exhibit characteristic oxidation states:
- Group 1 (Alkali Metals): The oxidation state of these elements is usually +1.
- Group 2 (Alkaline earth metals): Here, the typical oxidation state is +2.
- Group 17 (Halogens): These often display an oxidation state of -1. However, their oxidation state can become positive when combined with oxygen or other halogens.
- Group 18 (Noble Gases): These elements have oxidation state of 0 because they are mostly inert and do not form compounds easily.
Oxidation states in a period
As you move from left to right across a period in the periodic table, oxidation states generally become more positive due to the increasing nuclear charge. Elements on the right side of a period gain electrons to achieve a full outer shell, which usually leads to negative oxidation states. Here is a simple representation:
Oxidation states one group down
As you move down a group in the periodic table, the oxidation states of the elements generally remain the same. This is because each group has the same number of electrons in its outermost shell, which mainly determines their chemical reactivity and oxidation states.
Special cases
While general rules apply widely, some elements exhibit different oxidation states in different compounds due to special conditions or external factors. For example, transition metals often exhibit multiple oxidation states. Consider chromium and manganese, where their normal oxidation states can vary widely within a series of compounds:
Cr
: usually displays +2, +3, and +6Mn
: can show +2, +4, +6, and even +7
The role of oxidation states in redox reactions
Redox reactions are the main processes of oxidation and reduction, which involve the transfer of electrons between chemical species. Understanding oxidation states is important for balancing redox reactions, because a change in oxidation state is always complemented by an equivalent change in the opposite direction. This means that the total increase in oxidation states must equal the total decrease.
Here's a simple example of a redox reaction:
Example: Oxidation of zinc with copper sulphate:
Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s) Zn oxidation: 0 to +2 (loss of 2 electrons) Cu reduction: +2 to 0 (gain of 2 electrons)
Conclusion
The oxidation state is an important concept in chemistry that provides crucial insight into the behavior of elements during chemical reactions. The method of assigning oxidation states provides a structured way to understand electron transfer processes and makes balancing chemical equations more intuitive. By understanding oxidation states, chemists can predict reaction products, understand the electronic nature of bonds, and explore the rich variety of chemical reactions. With continued study, the application of oxidation states extends beyond basic chemistry, playing a fundamental role in advanced studies in inorganic, organic, and physical chemistry.