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The terminal C atom in butane is _______ hybridized.
(A) $ds{{p}^{2}}$
(B) $sp$
(C) $s{{p}^{2}}$
(D) $s{{p}^{3}}$

Last updated date: 13th Jun 2024
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Hint: There are 2 terminal carbons in the butane chain and both of them have the same hybridization. Both the terminal carbon atoms form 4 single bonds.

Complete step by step solution:
Butane is a member of the alkane group. It has 4 carbons. It has a straight-chain structure.
The formula of butane is ${{C}_{4}}{{H}_{8}}$
Since it is a member of the alkane group, all the bonds in butane will be a single bond.
So the structure of butane is given below:

So, there are two terminal carbon atoms in the butane, carbon number one and four.
Both the carbon atoms are connected to three hydrogen atoms and one propyl group.
There are some steps which can help to calculate the hybridization of an atom:
First, look at the atom and count the number of atoms or molecules to which it is connected. If the atom has lone pairs, then it is also counted.
Add both numbers and:
If the number is 4 then it is $s{{p}^{3}}$hybridized.
If the number is 3 then it is $s{{p}^{2}}$hybridized
If the number is 2 then it is $sp$hybridized
So, in both the terminal carbon atoms of butane the number is 4 (three hydrogen atoms and one propyl molecule).
Hence, the hybridization is $s{{p}^{3}}$.

So, the correct answer is an option (d)- $s{{p}^{3}}$.

Note: $s{{p}^{2}}$hybridization mostly occurs when the carbon atom has a double bond and $sp$hybridization occurs mostly when the carbon atom has a triple bond. $ds{{p}^{2}}$hybridization is shown by heavier elements because they have vacant d-orbitals.