bonding types

8/3/2019 Bonding Types

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Bonding Types

Given the electronic configurations of any element or ion, it is possible to determine thecombination of orbitals and suborbitals used in forming chemical bonds. This can be done on two

levels: (a) the valence bond atomic orbital method; and, (b) the molecular orbital method. The

former is used for descriptions of VSEPR geometry of simple covalent inorganic molecules; the

latter is used for descriptions of geometry in simple covalent organic molecules. These two

approaches can be combined to serve as more complete descriptors.

There are no orbital descriptions for ionic molecules, whether inorganic or organic, because

ionic bonds consist of electron transfer, not sharing, and are in principle existent only by virtue of

Coulomb's Law of Electromagnetic Attraction:

F = 1 q1q2

4 πε0 r2

F is the coulombic force of attraction or repulsion measured in Newtons, π has its usual

meaning, ε0 is the permittivity constant = 8.85 x 10-12 coulomb2 /Newton-meter2, q is the charge

on the first ion in coulombs, q is the charge on the second ion in coulombs, and r is the distance

between them in meters. Note: A charge of +1 or –1 on an ion equals 1.60 x 10-19 coulombs.

A permittivity constant is a measure of the ability for electromagnetic attraction to pass through

a particular medium, in this case air. It would have a different value for water, for example.

Larger integer ionic charges multiply the coulombic charge accordingly. The first term is a

constant, so the expression can be rewritten as:

F = (8.99 x 109

newton-meter2 /coulomb

2) q1q2

r2

Valence Bond Atomic Orbitals Given the electronic configuration of any element, only electrons in the valence orbital, the

outermost orbital, are involved in covalent bond formation. Valence orbitals include any empty or

partially filled suborbitals in that orbital. This means that nonionic elements in the third or later

period of the Periodic Table of Elements contain unfilled d or f suborbitals as part of their

valence orbitals.

Examples:

7N = 1s22s

22p

3is complete electron configuration; 2s

22p

3is incompletely filled valence orbital

15P = 1s22s22p63s23p33d0 is complete electron configuration; 3s23p33d0 is incompletely filled

valence orbital

Note: Pictures of s, p, d, f suborbitals are provided in the Quantum Numbers review.

Hybridization is the physical/mathematical process of mixing the energies of valence orbitals

to make them all energetically equivalent . Since s suborbitals are lower energy, more stable, than

p suborbitals, which are in turn lower energy than d suborbitals, there must be some benefit

gained through hybridization that outweighs making some suborbitals less stable. This benefit

arises from the uncoupling of paired electrons within a single suborbital. By separating electrons,

by placing them in different energetically equivalent suborbitals, the Coulombic repulsive force

between them is reduced .



When atoms form single bonds with other atoms, they hybridize all the completely or partially

filled suborbitals. When atoms form double bonds with other atoms, they hybridize all but one of

the completely or partially filled suborbitals. When atoms form triple bonds with other atoms,

they hybridize all but two of the completely or partially filled suborbitals . The unhybridized suborbitals in double or triple bonds are always p suborbitals.

Examples:

7N = 2s22p3 gains no benefit through hybridization because all valence orbitals are already

partially or completely filled. No uncoupling of paired electrons is possible. Although we

write the hybridization of Nitrogen in NH 3 as sp3 , this is a fiction , and employed only to

insert consistency in designations for the VSEPR theory approach.

15P = 3s23p33d0 gains substantial benefit through hybridization because the empty d suborbitals

permit the uncoupling of electrons in the 3s suborbital . This means Phosphorous

hybridizes to a configuration of 3s13p

33d

1, which is more simply written as sp

3d in

VSEPR theory for molecules such as PCl5.

The valence bond atomic orbital description of bonding in molecules simply lists the

contributions of suborbitals, hydridized or not, coming from each atom participating in covalent

bond formation.

Examples:(1) Each of the three N—H bonds in NH3 is described as sp3—s coming from sp3 hydbridized

suborbitals for Nitrogen and unhybridized s suborbitals from the three Hydrogens.

(2) Each of the five P—Cl bonds in PCl5 is described as sp3d—p coming from sp

3d hydbridized

suborbitals for Phosphorous and unhybridized p suborbitals from the five Chlorines.

Octet Rule violation for halogens (Cl, Br, I) when bonded to multiple oxygen atoms derives

from the rehybridization of the halogen atoms by use of empty d suborbitals.

Example:

17Cl = 1s22s

22p

63s

23p

53d

0rehybridizes its valence suborbitals to 3s

13p

33d

3thus permitting

formation of seven bonds to as many as four Oxygens for a total of 14 valence orbital

electrons.

The valence bond designation of HClO4 is s—sp3 for the H—O bond, sp3—sp3d3 for the

(H)O—Cl bond, and sp3d3—sp2 for the three Cl=O bonds:

O||

H—O—Cl=O

||

O

Molecular OrbitalsMolecular orbitals derive from the physical/mathematical mixing of atomic suborbitals

involved in covalent bond formation. This mixing of atomic suborbitals always produces pairs of

molecular orbitals, which are designated as bonding and antibonding orbitals. As the names

imply, the filling of bonding orbitals with electrons is a favorable process for molecular stability,

while the filling of antibonding orbitals with electrons is a molecule destabilizing process. For



this reason, bonding orbitals are always filled before antibonding orbitals with the available

electrons.

There are three types of molecular orbitals: (a) sigma, (b) pi, and (c) nonbonding.

Sigma bonding orbitals (σ): the most stable, lowest in energy, orbital; always the firstformed between two atoms, and always possesses s suborbital atomic character (σ means s)

Pi bonding orbitals (π): next most stable; always form the second and third bonds

between two atoms, and always possess unhybridized p suborbital atomic character (π means p)

Non-bonding orbitals (n): always consist of non-reactive lone pair electrons on such

elements as N, O, P, S, and halogens

Sigma antibonding orbitals (σ*): the least stable, highest in energy, orbital; is always the

last filled between two atomsPi antibonding orbitals (π*): between non-bonding and sigma antibonding orbitals in

energy

Examples: CH4 has four sigma single bonds formed from one sp3

hybridized Carbon suborbitals

and four s unhybridized Hydrogen suborbitals. The eight electrons fill all the bondingorbitals and leave all the antibonding orbitals empty.

___ ___ ___ ___ σ*

/ \

/ \

Carbon sp3 ↑ ↑ ↑ ↑ ↑ ↑ ↑ ↑ s 4 Hydrogen

\ /

\ /

↑↓ ↑↓ ↑↓ ↑↓ σ

C2H4 has four sigma single bonds formed from two sp2 hybridized Carbon suborbitals

and four s unhybridized Hydrogen suborbitals. It also has a sigma bond formed from two

sp2 hybridized Carbon suborbitals, and a pi bond formed from two p unhybridized carbon

suborbitals. The twelve electrons fill all the bonding orbitals and leave all the antibonding

orbitals empty.

___ ___ ___ ___ σ*C-H

__ σ*C-C

/ \

/ \ / __π*C-C \

/ \ p ↑

2 Carbon sp2 ↑ ↑ ↑ ↑ ↑ ↑ ↑ s 4 Hydrogen

\ /

\ ↑↓ πC-C /

\ /

\ /

↑↓ σC-C

↑↓ ↑↓ ↑↓ ↑↓ σC-H



Molecular Orbitals

bonding types

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