CHEMICAL BONDING

CHEMICAL BONDING

Chemical Bonding
There are numerous types of chemical bonds and forces acting jointly to combine molecules together. The two most fundamental types of bonds are ionic and covalent bond. In ionic bonding, atoms transfer electrons to each other. Ionic bonds need at least one electron donor and one electron acceptor. On the contrary, atoms that have similar electro negativity share electrons through covalent bonds as for such atoms, donating or receiving electrons are not favorable.

Chemical bonding is a means through which atoms unite to form molecules. Chemical bond exists between two atoms or groups of atoms when the forces acting between them are physically powerful enough to result to the formation of an aggregate with adequate stability to be termed an autonomous species. The no of bonds an atom forms matches up to the number of electron at its outer shell. Bond energy is the quantity of energy necessary to break a bond and create neutral atoms. In line with Coulomb's law every bond as a result of attraction that exist between unlike charges. On the other hand, the manner this force is manifested varies depending on the atoms concerned. The main types of chemical bond are the ionic, covalent, metallic, and hydrogen bonds. The ionic and covalent bonds are ideal forms but the majority of the bond types are of an intermediary type.

Bonding energy between two atoms
The interaction energy between two atoms at equilibrium is referred to as the bonding energy between the two atoms. To break the bond, this energy must be supplied from outside. Breaking the bond means that the two atoms become infinitely separated. In real substances that are made up of varieties of atoms, bonding is calculated by stating the bonding energy of the entire substances in terms of the disjointing distances among all atoms. There are different types of bonding:

• Primary bonding: Ionic (involves transfer of outermost electrons)

• Covalent (involves sharing of outermost electrons, directional)

• Metallic (involves delocalization of valence electrons)

• Secondary or van der Waals Bonding:(widespread, but less strong than primary bonding)

• Dipole-dipole

• H-bonds

• Polar molecule-induced dipole

• Variable dipole (the most weak bond)

The Ionic Bonding
Ionic bonding is the total transfer of outermost electron(s) between atoms. It is the type of chemical bond that produces two oppositely charged ions. In ionic bonds, the metal loses electrons to turn into a positively charged cation, while the non-metal receives those electrons to turn into a negatively charged anion. For ionic bond to occur there must be an electron donor, metal, and an electron acceptor, non metal.

Ionic Bonding is occurs because metals have a small number of electrons in their outmost orbital. Through the loss of those electrons, these metals can attain noble-gas configuration and meet the octet rule. Likewise, non metals that have nearly 8 electrons in their outermost shell have the tendency of readily accepting electrons to attain their noble gas configuration. In ionic bonding, over 1 electron can be donated or received to fulfill the octet rule. The charge on the anion and cation matches up with the number of electrons contributed or received. In ionic bonds, the net charge of the compound must be zero.

The ionic bond is a chemical bond formed as a result of attraction between two opposite charged ions. The atoms of metallic elements like sodium easily lose their valence electrons, whereas the atoms of non-metals like chlorine have the tendency to gain electrons. The reaction between them results to a highly stable ions which maintain their individual structures while approaching one another to form a stable molecule or crystal. In an ionic crystals such as sodium chloride, no separate diatomic molecules are present; instead, the crystal is made up of composed of independent Na+ and Cl− ions, with each being attracted to adjoining ions of the opposite charge giving rise to one single gigantic molecule.

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The covalent bond
The covalent bond is the type of bond formed when two atoms share a pair of electrons. There is no net charge on each of the atoms; the attractive force between them is formed through the interaction of the pair of electron within the nuclei of both atoms. When the interacting atoms share more than two electrons, it leads to the creation of double and triple bonds. This is because every shared pair forms its own particular bond. The interacting atoms, by sharing their electrons are able to attain a highly stable electron configuration that correspond to that an inert gas.

Covalent Bonding
Covalent bonding is the chemical bond that results from sharing of electrons between atoms. This type of bonding arises between two atoms of the identical element or elements that are placed close to each other in the periodic table. Covalent bonding occurs mainly between non metals; but it can as well be found between non metals and metals alike.

Polar and non-polar physical properties of covalent compounds

Physical properties Covalent compounds
States (at room temperature) Solid, liquid, gas
Electrical conductivity Usually none
Electrical conductivity Usually none
Boiling point and Melting Varies, but usually lower than ionic compounds
Solubility in water Varies, but usually lower than ionic compounds
If we take methane (CH4), as an example; in methane, carbon shares one electron pair with each of the 4 hydrogen atom; so that the total number of electrons shared by carbon is eight, which matches up with the numbers of electron in the valence shell of neon; every hydrogen atom shares two electrons, which matches up with the electron configuration of helium.

In the majority of covalent bonds, every one of the atom contribute just a single electron to the shared pair. In some instances however, the two electrons are donated from the same atom. When this happens, it gives rise to a partially ionic character giving rise to what is called a coordinate link. In actual fact, a purely covalent bond can only be found among two identical atoms.

Covalent bonds play a significant role in organic chemistry due to the capability of the carbon atom to form 4 covalent bonds. These bonds are arranged in specific directions in space, resulting to the complex geometry of organic molecules. If every one of the four bonds is one just like in methane, the resulting molecule will have the shape of a tetrahedron. The significance of shared electron pairs was originally discovered by the

American chemist G. N. Lewis (1916), who stated that there are only extremely few stable molecules that exist which have a total numbers of electrons that is odd. His octet rule permits chemists to forecast the most likely bond structure and charge allocation for molecules and ions. With the initiation of quantum mechanics, it was recognized that the electrons in a shared pair must possess opposite spin, as required by the Pauli Exclusion Principle.

The molecular orbital theory
The molecular orbital theory was developed to foretell the accurate sharing of the electron density in a variety of molecular structures. The American chemist Linus Pauling initiated the concept of resonance to clarify how stability is achieved when above one practical molecular structure is achievable: the actual molecule is a coherent mixture of the two structures.

Metallic and hydrogen bond
Contrary to ionic and covalent bonds, which are copiously available in a great number of molecules, the metallic and hydrogen bonds are highly specific.

The metallic bond is the bond accountable for the crystalline structure of pure metals. This bond cannot be ionic due to the fact that every one of its atoms is identical. It cannot also be covalent ordinarily due to the fact that there are very few valence pairs of electrons to be shared among adjoining atoms. As an alternative, the valence electrons are shared jointly by all the atoms in the crystal. The electrons act as a free gas moving within the lattice of preset, positive ionic cores. The intense mobility of the electrons in a metal gives explanations for its high thermal and electrical conductivity.

Hydrogen bonding is a very powerful electrostatic attraction between two independent polar molecules. That is between the molecules in which the charges are unequally dispersed, characteristically containing nitrogen, oxygen, or fluorine. These elements have physically powerful electron-attracting power, and the hydrogen atom serves as an overpass between them. The hydrogen bond, which plays a crucial role in molecular biology, is by far weaker than the ionic or covalent bonds. Hydrogen bond is answerable for the structure of ice.

Bonding in Organic Chemistry
Ionic and Covalent bonds are the two top limits of bonding. Polar covalent is the midway type of bonding in between the two extremes. Some ionic bonds possess covalent character and some covalent bonds are partly ionic. For instance, the majority of carbon-based compounds are covalently bonded, however they can also be partly ionic.

Polarity is an evaluation of the separation of charge in a compound. A compound's polarity is reliant on the symmetry of the compound together with the differences in electronegativity between atoms. Polarity arises when the electron pushing elements from the left side of the periodic table, shares electrons with the electron-pulling-elements from the right side of the period table. This generates a spectrum of polarity, with ionic (polar) at one end, covalent (non-polar) at the other end, and polar covalent in the center.

In co-operation, these bonds are vital in Organic Chemistry. Ionic bonds are significant because they allow the synthesis of specific organic compounds. Covalent bonds are particularly essential since most carbon molecules interact primarily through covalent bonding. Covalent bonding permits molecules to share electrons with other molecules, forming long chains of compounds and giving rise to more complexity in life.

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Polar and Non-Polar Shapes
Molecules that have a linear, trigonal planar, tetrahedral, trigonal bipyramidal, or octahedral shape, are non-polar in nature. These are shapes which do not have non-bonding lone pairs like Methane, CH4. However if a few bonds are polar while the rest are not, there will be a general dipole, and the molecule will be polar. Example Chloroform, CHCl3.

Dipole-Dipole Bonds
When two polar molecules come close to each other, they will position themselves in order to let the negative and positive sides' line up. There will be an attractive force linking the two molecules together, but it is not virtually as well-built a force as the intramolecular bonds. This explains the way different types of molecules bond together to form bulky solids or liquids.

Van der Waals forces are initiated by temporary dipoles formed when electron locations are asymmetrical. The electrons are continually tracking the nucleus, and by chance they could come very close together. The uneven concentration of electrons could result to one side of the atom becoming more negatively-charged than the other, forming a temporary dipole. Van der Waals forces are the reason why nitrogen can be liquified.