5 Types of Bonds in Chemistry with Diagrams
There are different chemical species, like atoms, ions, and molecules. A chemical bond is a lasting attraction between two chemical species. Depending on the specific type of bond and the nature of the particular species, the bond may be strong or weak.
Chemical species usually bond to fill their electron orbitals and shells, due to the attraction between two oppositely charged ions, due to the electronegativity difference between two atoms in separate molecules, etc.
The amount of energy required to break one mole of a specific bond is known as the bond enthalpy or bond energy of that bond.
All bonds can be classified into two types –
- Primary (strong) bonds- These include the bonds which are strong and difficult to break like covalent, ionic and metallic bonds.
- Secondary (weak) bonds – This class includes the weaker more temporary bonds like London dispersion forces, dipole-dipole force and hydrogen bonding.
This type of bond is also called the molecular bond. In this bond, the electrons are shared among the two atoms. Thus each covalent bond is made of a pair of electrons, one electron belonging to each chemical species. This type of bond is formed usually between two nonmetal atoms with similar or almost the same electronegativity values. This sharing of electrons usually occurs such that the atoms involved can fill their outer orbitals to achieve a stable noble gas configuration. The electron pair that take part in the covalent bond is called a bonding pair/ shared pair. The electron pairs not taking part in the bond are known as lone pairs. The bond pair and the bond is represented as a straight line between the two atoms. Each lone pair is represented as two dots on top of the atom that it belongs to.
Covalent bonds have interactions and overlap of the sigma and pi orbitals in the two atoms. Thus we can have a covalent single, double and triple bond.
Covalent bonds are of two types as
- Nonpolar covalent bond and
- polar covalent bond.
A non-polar bond is one where the two atoms of the molecule share the atoms equally, it is known as non-polar/pure covalent bond. This occurs when the atoms participating in the bond are the same. This is the case in hydrogen gas ( H2), oxygen gas( O2) nitrogen gas (N2), etc.
A polar covalent bond is present in compounds or radicals where the two atoms are of varying electronegativities like NH3 or CH4. In this case, the electrons are pulled more towards one atom ( the more electronegative one) than the other. Thus there is unequal sharing. More the electronegativity difference between the two atoms, the more polar the bond is.
The ionic bond is formed due to an electrostatic force of attraction between two oppositely charged ions. This bond is usually formed between a metal and a nonmetal. Prior to the electrostatic attraction, there is a total transfer of valence electrons from one atom to the other atom.
The first atom (usually a metal which is electropositive in nature) loses electrons to achieve closest noble gas configuration, fulfill the octet rule or have completely filled outer orbitals. This loss of electrons is facilitated by ionization energy and makes the ion stable. The second atom (usually a nonmetal which is electronegative in nature) accepts electrons to get closest noble gas configuration, fulfill the octet rule or have completely filled outer orbitals. This is facilitated by electron gain enthalpy. The oppositely charged ions get attracted due to electrostatic forces and bond forms between them, thus releasing energy.
Ionic compounds dissociate into ions in water. A common example of ionic compounds is NaCl, CaCl2, etc.
Metals have low ionization energy. Thus they release electrons easily as they are electropositive. These electrons leave their parent atoms and get delocalized. Metal bonding occurs among metal atoms.
For example, in a bar of sodium metal. Sodium has an electronic configuration of 1s2 2s2 2p6 3s1.
The 3s orbital is the outermost orbital and in the metal bar when all the atoms are close together the 3s orbital of one atom overlaps with the 3s orbital of the neighboring atom. Thus the atomic orbitals overlap to form molecular orbitals. This overlap occurs with all the neighboring atoms on each side. Thus a huge number of molecular orbitals extend throughout the bar. The electrons leave the 3s orbital and travel through the molecular orbitals thus getting delocalized. The metal is thus held together by the strong forces of attraction between the positive nuclei and delocalized electrons. Thus metallic bonding is very strong and metals have high melting and boiling points. Metallic bonding can also be found in copper, aluminum, etc.
This occurs when hydrogen is directly bonded with a highly electronegative element like oxygen or fluorine. In this case, the hydrogen acquires a partial positive charge because the electrons in the bond get more attracted to the electronegative atom. If there is another electronegative atom nearby, the hydrogen atom gets attracted to the electronegative atom and forms a hydrogen bond.
Stronger the electronegativity difference, stronger is the bond. The hydrogen actually gets attracted to the lone pairs of the electronegative atom. This type of bonding can be found in HF, H2O, etc.
In a covalent bond if there is electronegativity difference between the two atoms, then the electrons get pulled more towards the electronegative atom and it has higher electron density. This results in a molecular dipole. The dipole moment is from the direction of the electropositive to the electronegative atom. When the dipole moments are not canceled due to geometry of the atom it becomes a polar atom. In each molecule, one atom is partially positive and the other is partially negative. The attractive forces between the partial positive of one atom and the partial negative of another atom are known as dipole-dipole interactions. These forces can be seen in HCl.