What is a bond order of fluorine

The chemistry side

Subsections

Atomic bonds

Lewis formulas

Atoms have an outer electron shell that Valence electron shell is called. This is where the electrons are located, which are important in chemical reactions. These electrons are now drawn in Lewis formulas using a point for single electrons, one line for an electron pair. Lewis hypothesized that atoms bonded to one another share electrons. You can draw a pair of electrons between two atoms as a sign that it is a bonding pair of electrons. So you get the Lewis formulas also Valence stroke formulas called.

Octet rule

Atoms try to divide their electrons so that they 8 electrons if possible in the outer shell. This is imperative for the elements fluorine, oxygen, nitrogen and carbon in the first eighth period. These elements have as Binding order the Number of electronswho give them absence.

Exceptions to the octet rule

Elements that already have a noble gas configuration cannot (mostly) form bonds with other elements.
Lithium can hold 1 electron. Beryllium 2, because a p-orib valley is occupied. Boron can take up 3 electrons because two p orbitals are occupied. The cohesion is 1,2 and 3 respectively.

Atoms that come from higher periods (such as sulfur or phosphorus can also simply fill d-orbitals with electrons, which is why they can have more than 8 electrons in the valence electron shell.

Determination of the bond order using the Lewis formulas

To meet the octet rule, some atoms have to form multiple bonds with other atoms. So has for example two double bonds, otherwise neither the carbon atom nor the two oxygen atoms fulfill the octet rule. even has a triple bond. A free electron pair of the oxygen atom is still used as a bond pair, since otherwise carbon does not fulfill the octet rule (dative bond). So one can determine the binding order with the help of the four elements of the first eighth period, which strictly fulfill the octet rule.

Formal charges

Formal charges are purely fictitious charges and have practically no meaning. Charges in a molecule (except for ions) are not located on one side of the molecule, but electrons are only slightly "pulled" in one direction (see Partial Charges).
The formal charge is determined by simply cutting all bonds in the middle so that one electron from each bonded pair of electrons goes to one atom and the other to the other atom. Then the electrons are counted. If the atom has more electrons than in the original state, it is negative. Has it been less positive.
One writes

It can happen that formal charges in a molecule are not balanced again. This molecule is then an ion such as the carbonate ion or the -Ion.

Partial charges

Partial charges are actual charges. If an atom of the molecule, because of its higher electronegativity, attracts the binding electrons more than other atoms, it is negatively charged and the other atoms are positive. One writes:

In water, for example, the oxygen atom is partially more negatively charged than the hydrogen atoms.

Mesomeric limit formulas

There are molecules and ions in which double bonds and lone pairs of electrons can be attached in different ways. These molecules have a mesomerism. Not one possibility of the arrangement of the electron pairs is chosen, but all at the same time, whereby all possibilities are blurred within one another and the formulas drawn with the Lewis formulas above are only limiting formulas, which almost never occur.
Have different mesomeric limit formulas
  • the same geometry
  • comparable energy of the structures
  • the same sum of lone pairs of electrons, bonded electron pairs and the same sum of formal charges
  • always negative partial charges on the most electronegative elements.
In the case of mesomeric limit formulas, the electron pairs between the drawn limit structures must not jump over an intermediate state without drawing it.

VB theory (Valence Bond theory)

The experimentally determined binding energy of amounts to . The VB theory describes how this energy comes about.
A, B: atomic orbitals
1,2: electrons

LCAO method (linear combination of atomic orbitals)

In order to approximate the molecular orbitals, two atomic orbitals are added as a binding molecular orbital or subtracted as an antibonding molecular orbital in the LCAO method. Let us consider the molecular orbitals of the two molecules A and B:


To the overall function of all molecular orbitals, all occupied molecular orbitals are multiplied together. With molecular hydrogen there are two binding molecular orbitals, each occupied by an electron:


Different molecular orbitals

Molecular orbitals are thus formed by overlapping atomic orbitals. There are s, p, d etc. atomic orbitals. Only positive orbital parts that overlap can form a bond. If a positive part equally overlaps a positive part and a negative part of another atomic orbital, there is no bond. There are two different types of bindings:

σ bondThe atomic orbitals overlap by placing one orbital on the other at the end. A σ-bond exists, for example, when a p-orbital is added to an s-orbital by placing the p-orbital across and overlapping the s-orbital at one end. The notation for the orbital just introduced is σsp. There is exactly the same σ2p-Molecular orbitals.
π bondThis bond is created when two atomic orbitals are placed alongside one another. For example, a p orbital can longitudinally overlap another p orbital by overlapping the positive part with the positive part and the negative part overlapping the negative part. Such an Orib valley is called π2p-Orbital.