Covalent Bond theory
Covalent Bond (nonmetal-nonmetal):
There are many molecules whose formation can not be explained by the ionic bond, the simplest cases are molecules which are composed of identical atoms, such as H2, Cl2, etc. In these cases, called homonuclear molecules as the two atoms have identical electronic configuration as they are equal, it is impossible that an atom yield an electron to the other, which is clearly seen, taking into account the ionization energies and electron affinities of atoms and doing an energy balance of the process of forming ions, which require a lot of energy.
In Homonuclear molecules the molecular orbital is located symmetrically between the two atoms and electrons, which are forming the bond, are truly shared by two atoms.
When the two atoms which are forming the molecule are different, heteronuclear molecules, electrons which form the bond are not equally shared, but on average, are closer to one core than the other. If we take for example the molecule of hydrogen chloride, the chlorine atom has more appetite for electrons than the hydrogen atom, which is measured by the electron affinity or electronegativity which is greater in the case of chlorine in hydrogen. Therefore, in the formation of the molecule of hydrogen chloride, chlorine strip more electrons, which the electronic distribution is not symmetric. Thus, a separation between the center of positive charge and the center of negative charges with an electric dipole moment is created appears. It is then said that the molecule is polar, or in polyatomic molecules, the bond is polar. This can also be expressed by saying that the link is noncovalently pure, but has some ionic character, which depends primarily on the difference in electronegativity between the two atoms forming the linkage.
For Chlorine Oxygen or water molecule, the oxygen atom is surrounded by eight electrons, counting twice the link electron pair. By achieving these eight electrons, the atom O satisfies the Octet Rule, requirement to have eight electrons in the valence shell of each atom in a Lewis structure, but the H atom is an exception because it can only has two electrons in the valence shell.
Sharing a single pair of electrons between bonded atoms gives rise to a single covalent bond.
Therefore, the links can form an element with different others, there is a continuous degradation, ranging from pure covalent bond to almost pure ion passing through many intermediate states. In diatomic molecules, the ionic bond character is directly related to the dipole moment of the molecule. In polyatomic molecules things are not so simple, but can also relate the ionic character of the bonds with what are called dipole moments of connection, resulting from decomposing the dipole moment of the molecule in a vector sum of partial moments, associated each one to one link.
Coordinate Covalent Bond:
It may be the case that one of the two electrons contribution atoms while the other input a vacant orbital only. In this case the link is called coordinated or dative bond donor and acceptor, the donor or donor being the atom that provides the electron pair acceptor atom and which contributes only the vacant orbital.
Once the link has been formed, you can not determine which of the links is the coordinate covalent, being almost impossible to distinguish between normal and coordinate covalent bond covalent bond.
Multiple Covalent link:
Often it is needed to share more than one pair of electrons to achieve octet (electronic configuration of noble gas). This is what happens for example in the molecule of carbon dioxide, where one C atom can share a valence electron with an O atom, forming two carbon-oxygen single bonds.
This does not make a byte or C atom or two of O. The problem is solved by moving unpaired electrons towards the region of the link. Linked together two atoms share electron pairs forming a covalent double bond.
In the case of the N2 molecule, the bond is covalent triple. Double and triple covalent bonds are called multiple covalent bonds.
Covalent polar link:
Chemical Bonds have been introduced as belonging to two clearly distinct categories: ionic bonds with electrons transferred completely and covalent bonds with pairs of electrons shared equally.
A covalent bond in which two atoms not equally share electrons is called polar covalent.
In the case of H2 and Cl2 atoms, the two atoms have the same electron affinity and electron moving toward either. The centers of positive and negative charge coincide: both are at a point equidistant from the two atomic nuclei. The links H-H and Cl-Cl are nonpolar.
In HCl, the chlorine atom attracts electrons more strongly than H, the electronic charge density is greater in the vicinity of Cl atom in the vicinity of the H atom negative charges center is closer to the chlorine core center positive charges. It is said that there is a charge separation in the H-Cl bond and the bond is polar. HCl The polar link can be represented by a Lewis structure in which the pair of nonbonding electron Cl is closest to that H.
Strategy for writing Lewis structures correctly:
- Determine the total number of valence electrons of the structure.
- Identify the central atom and the terminal atoms.
- Write an appropriate structural skeleton and joins the backbone atoms by single covalent bonds.
- For each link of the skeleton, subtract two electrons of the total number of valence electrons.
- electrons remaining valence full octets first terminal atoms and then complete, as far as possible octets central atom or atoms.
- missing octet to one or more central atoms, electrons move lone pairs of terminals atoms forming multiple covalent bonds with the central atoms.
They are apparent charges appearing on some atoms Lewis structure when atoms have not contributed the same number of electrons to the covalent bond that unites them.
The formal charge of an atom in a Lewis structure is the number of valence electrons in the free atom (not combined) minus the number of electrons assigned to that atom in the Lewis structure.
In a Lewis structure electrons are assigned to atoms as follows:
- Count all lone pairs of electrons as belonging entirely to the atom that is.
- Divide all bonding pairs of electrons equally between the bonded atoms.
General rules that help us to determine if a Lewis structure is acceptable based on their formal charges:
- The sum of the formal charges of atoms in a Lewis structure must be zero for a neutral and equal to the charge for a polyatomic molecule ion.
- If are necessary formal charges, they must be as small as possible.
- Negative formal charges usually appear on more electronegative atoms and positive formal charges in less electronegative atoms.
- The structures with formal charges of the same sign on adjacent atoms are unlikely.
You can download the App BioProfe READER to practice this theory with self-corrected exercises.