Covalent Bond
A covalent bond is a bond formed when two atoms share a pair of electrons. A neutral group of atoms bonded together covalently is called a molecule, and a substance which is made up of molecules is called a molecular substance. Covalent molecules make up many common substances, including plastics, paper, and human tissue.
Another type of bond that occurs when electrons are transferred between atoms is called an ionic bond(or electrovalent bond). Compounds held together by ionic bonds are called ionic compounds. Molecular compounds have lower melting and boiling points than do ionic compounds. Molecular compounds often occur as individual molecules, whereas ionic compounds occur as a crystal lattice. This crystal lattice structure gives the ionic compounds their higher melting and boiling points.
One characteristic that both ionic and covalent compounds share is that they both adhere to the octet rule. The octet rule is the principle that describes the bonding in atoms. Individual atoms are unstable unless they have an octet of electrons in their highest energy level. The electrons in this level are called valence electrons. When atoms gain, lose, or share electrons with other atoms, they satisfy the octet rule and form chemical compounds. An example of this is the formation of molecular fluorine (F2) from two individual fluorine atoms. A fluorine atom has seven valence electrons. According to the octet rule, eight valence electrons is the most stable configuration, so each fluorine atom needs one more electron. If each fluorine atom shares one electron with the other, the octet rule will be satisfied and a covalent bond is formed. Another example of a molecule formed by covalent bonds is ammonia (NH3). A nitrogen atom has five valence electrons, so it needs three more to satisfy the octet rule.Hydrogen has one valence electron, so it needs one more to fill its outer energy level (hydrogen has electrons in only the first energy level, which fills when two electrons are present, unlike the other energy levels that require eight). If one nitrogen atom shares an electron with each of three hydrogen atoms, and the three hydrogen atoms each share an electron with the nitrogen atom, three covalent bonds are formed and the octet rule is satisfied for all four atoms.
Covalent bonds are formed due to the forces of electric attraction between atoms, for example, the covalent bond that forms between two fluorine atoms. An attractive force occurs between the negatively charged electrons on the first fluorine atom and the positively charged protons on the second fluorine atom. A repulsive force also occurs between the electrons in each atom and between the protons in each atom. It would seem that this repulsive force would cause the atoms to move away from each other. The distance between the protons in one fluorine atom and the protons in the other fluorine atom (and, similarly, between the electrons between each atom as well) is greater than the distance between the electrons of one atom and the protons of the other. Therefore, the attractions between the two atoms are greater than the repulsions, and the two fluorine atoms are held together. Making the individual flourine atoms more stable (satisfying the octet rule) is also important. The stabilization of the atoms outweighs any repulsive forces there may be between the two atoms.
The fact that there are repulsive as well as attractive forces between two atoms means that the covalent bond must be flexible. If two atoms start to move apart, the attractive forces will draw them back toward each other. If they then move too close to each other, the repulsive forces will push them apart. The atoms in a covalent bond are actually vibrating back and forth around an average distance where the two forces are balanced. This average distance is called the bond length. The bond length is related to bond energy, which is the energy required to break a covalent bond. A general rule relating bond length and bond energy is the shorter the bond length, the greater the bond energy; the two are inversely related.
Covalent bonds can occur as single bonds or multiple bonds. Molecular fluorine and ammonia are both examples of single covalent bonds--in each case, a single pair of electrons is shared between atoms. In a multiple bond, two atoms share more than one pair of electrons. For example, a double covalent bond occurs when two atoms share two pairs of electrons. This occurs in a formaldehyde (H2CO) molecule.Carbon has four valence electrons, so it needs four more to satisfy the octet rule. It can only share one pair of electrons from each of the hydrogen atoms for a total of two valence electrons gained. In order to satisfy the octet rule and gain two more valence electrons, it must share two pairs with the oxygen atom. Oxygen has six valence electrons, and when it shares two pairs with the carbon atom it gains the two it needs to satisfy the octet rule. The two shared pairs of electrons between carbon and oxygen create a double covalent bond.
Two atoms can share three pairs of electrons as well, forming a triple covalent bond. This occurs in an ethyne (C2H2) molecule. Each carbon atom needs to gain four valence electrons. Each can gain one by sharing a pair of electrons with a hydrogen atom, and then gain the other three by sharing three pairs of electrons with each other. The three shared pairs of electrons between the two carbon atoms create a triple covalent bond.
When electrons are shared between atoms, they are not always shared equally. One atom may attract the electrons more than the other. This has to do with an atom's electronegativity. Electronegativity is a way to measure how much an atom attracts electrons in a chemical bond. The greater the electronegativity, the greater the attraction. When two atoms with two different electronegativities share electrons, the atom with the greater electronegativity will have a stronger attraction towards the electrons. When one atom has a much greater electronegativity than the other, the covalent bond between them is said to be polar. A polar bond results in an uneven distribution of charge. The atom with the greater electronegativity becomes slightly negative because it has a greater electron density, while the atom with the lower electronegativity becomes slightly positive. A molecule that has a polar covalent bond is called a polar molecule. An example of a polar molecule is water (H2O). Oxygen is more electronegative than hydrogen. The electrons that are shared between the oxygen and the two hydrogen atoms are pulled closer to the oxygen atom, giving the oxygen a slight negative charge. The hydrogen atoms, as a result, have a slight positive charge.
When two atoms with similar electronegativities share electrons, the covalent bond that results is said to be nonpolar. In a nonpolar covalent bond, both atoms attract the electrons equally. Nonpolar covalent bonds occur in covalent bonds between atoms of the same element, for example, molecular fluorine. Both atoms have the same electronegativity, so they exert an equal pull on the electrons shared between them. Electronegativity can be used to predict whether a bond will be a nonpolar covalent bond, a polar covalent bond, or an ionic bond. If the difference between the electronegativities of two atoms is 2.0 or greater, the bond is an ionic bond. If the difference is 0.4 or less, the bond is nonpolar. If the electronegativity difference between two atoms is between 0.4 and 2.0, the bond is considered to be a polar covalent bond. The greater the electronegativity difference, the greater the polarity.
Covalent bonds, or bonds formed by the sharing of electrons between two atoms, can occur in several forms. Single, double, and triple covalent bonds can occur. A covalent bond can also be polar or nonpolar. Covalent bonds between atoms are quite common and result in many different kinds of molecules. The phenomenon of two atoms sharing electrons is what builds many of the structures, both living and nonliving, in the world around us and understanding covalent bonds is an important part of both the physical and the life sciences.
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