At the atomic level, the bond order corresponds to the number of electron pairs of two atoms that are joined together. For example, the diatomic nitrogen molecule (N≡N) has a bond order of 3 because there are three chemical bonds joining the two atoms. According to the theory of molecular orbitals, the bond order is also defined as half the difference between the number of binding electrons and that of anti-binding electrons. To easily get the result you can use this formula:
Bond order = [(Number of electrons in a molecular bond) - (Number of electrons in a molecular antibonding)] / 2
Steps
Part 1 of 3: Quick Formula
Step 1. Learn the formula
According to the theory of molecular orbitals, the bond order is equal to the semidifference between the number of binding and antibonding electrons: Bond order = [(Number of electrons in a molecular bond) - (Number of electrons in a molecular antibonding)] / 2.
Step 2. Understand that the higher the bond order, the more stable the molecule will be
Each electron that enters a bonding molecular orbital helps stabilize the new molecule. Each electron that enters an antibonding molecular orbital destabilizes the molecule. Note that the new energy state corresponds to the bond order of the molecule.
If the bond order is zero, then the molecule cannot form. A very high bond order indicates greater stability for the new molecule
Step 3. Consider a simple example
Hydrogen atoms have one electron in the "s" orbital and this is capable of holding two electrons. When two hydrogen atoms bond together, each fills the "s" orbital of the other. In this way, two binding orbitals were formed. There are no other electrons that have been pushed to a higher energy level, the "p" orbital, so no antibonding orbitals have formed. In this case the bond order is (2−0) / 2 { displaystyle (2-0) / 2}
che è pari a 1. Questo genera la comune molecola H2: il gas idrogeno.
Parte 2 di 3: Visualizzare l'Ordine di Legame di Base
Step 1. Determine the binding order at a glance
A single covalent bond has a bond order of one, a covalent double bond corresponds to a bond order of two, a covalent triple bond has a bond order of three, and so on. In very simplistic terms, the bond order corresponds to the number of electron pairs holding two atoms together.
Step 2. Consider how atoms come together to form a molecule
In each molecule, the atoms are linked together by pairs of electrons. These revolve around the nucleus of a second atom of "orbitals" in which there can only be two electrons. If an orbital is not "full", that is, it has only one electron, or it is empty, then the unpaired electron can bond with the free electron of another atom.
- Depending on the size and complexity of a particular atom, it could have only one orbital or even four.
- When the nearest orbital is full, new electrons begin to gather in the next orbital, outside the nucleus, and continue until this "shell" is also complete. This process continues in increasingly larger shells, as large atoms have more electrons than small ones.
Step 3. Draw the Lewis structures
This is a very useful method for visualizing how the atoms in a molecule bond together. It represents each element with its chemical symbol (for example H for hydrogen, Cl for chlorine and so on). It represents the bonds between them with lines (- for the single bond, = for the double bond and ≡ for the triple bond). Identify the electrons not involved in the bonds and those coupled with points (for example: C:). Once you have written the Lewis structure, count the number of bonds and you will find the bond order.
The Lewis structure for the diatomic nitrogen molecule is N≡N. Each nitrogen atom has one pair of electrons and three unpaired electrons. When two nitrogen atoms meet, they share six unpaired electrons that intertwine in a powerful triple covalent bond
Part 3 of 3: Calculate the Bond Order according to Orbital Theory
Step 1. Consult a diagram of the orbital shells
Remember that each shell moves further and further away from the nucleus of the atom. Following the property of entropy, energy always tends to the minimum equilibrium state. So the electrons try to first occupy the available orbitals closest to the nucleus.
Step 2. Learn the difference between bonding and antibonding orbitals
When two atoms join together to form a molecule, they tend to use their respective atoms to fill the orbitals with the lowest energy level. The binding electrons are, in practice, those that come together and fall to the lowest energy level. Anti-binding electrons are the "free" or unpaired electrons that are pushed into an orbital with a higher energy level.
- Bonding Electrons: By looking at the number of electrons present in each atom's orbitals, you can determine how many electrons are in the higher energy state and which can fill a more stable shell with a lower energy level. These "fill electrons" are called binding electrons.
- Anti-bonding electrons: when two atoms unite to form a molecule they share some electrons, some of these are brought to a higher energy level, then to an outer shell as the inner ones and with a lower energy level fill up. These electrons are called antibonders.
