Intermolecular Bonding Essay

Topics: Chemical bonding, Atom, Chemical bond Pages: 5 (1620 words) Published: October 8, 1999
Intermolecular Bonding Essay

Write an essay on intermolecular bonding. Explain how each type of bond arises and the evidence for the existence of each. Comment on their strengths in relation to the types of atoms involved; the covalent bond and relative to each other. Use the concepts of different types and strengths of intermolecular bonds to explain the following:

There exists four types of intermolecular bonding, they include ionic, covalent, Van der waals and hydrogen bonding. In order to describe the existence of such bonding you must also understand the concepts of polarity, polar and non-polar, and electronegativity.

Ionic bonds are created by the complete transfer of electrons from one atom to another. In this process of electron transfer, each atom becomes a ion that is isoelectronic with the nearest noble gas., the substance is held together by electrostatic forces between the ions. The tendency for these ions to be formed by elements is corespondent to the octet rule, when atoms react,, they tend to do so in such a way that they attain an outer shell containing eight electrons. The factors that effect the formation of ions are ionization energy, electron affinity, lattice energy.

Figure 1

The transfer of electrons involved in the formation of (a) sodium chloride and (b) calcium fluoride. Each atom forms an ion with an outer shell containing eight electrons.

For many elements, compounds cannot be formed by the production of ions, since the energy released in the formation of the lattice of ions would be insufficient to overcome the energy required to form the ions would be insufficient to overcome the energy required to form the ions in the first place. In order for the atoms to achieve a noble gas configuration they must use another method of bonding by the process of electron sharing. From figure 2, you can see that the example of two hydrogen atoms combing. As the atoms get closer together, each electron experiences an attraction towards the two nuclei and the electron density shifts so that the most probable place to find the two electrons is between the two nuclei. Effectively each atom now has a share of both the electrons. The electron density between the two nuclei exerts an attractive force on each nucleus keeping them held tightly together in a covalent bond.

Figure 2

A covalent bond forming between two hydrogen atoms.

It is also possible for two atoms share more than one pair of electrons, sharing two pairs results in a double bond and sharing three pairs results in a triple bond. Electronegativity is a measure of how powerful a atom is in a molecule to attract electrons. Polarization is a term given to name the unequal sharing of electrons in a covalent bond. Molecules that have unequal sharing of electrons are called polar molecules and dipole molecules are ones which have the charge separated, therefore all polar molecules must have a dipole attraction. Non- polar molecules are ones in which there shapes are symmetrical so the electrons are evenly distributed. Polar molecules have a permanent dipole in other words they have a permanent separation of charge. As a result of this, polar molecules are attracted to one another by forces called permanent dipole-permanent dipole interactions, in which the negative end of one molecule is attracted towards the positive end of another. These interactions decrease quite rapidly as the distance between molecules increases. They are approximately 100 times weaker than covalent bonds. There are also very strong types of dipole-dipole interactions called Hydrogen bonds. Evidence for the existence of such intermolecular forces lies in the properties of hydrides formed by element in groups 4,5,6 and 7. While all the hydrides formed in group 5 behave in a similar way, the hydrides of other groups do not. This suggest that the intermolecular forces in these hydrides are much stronger than expected compared with other hydrides of the...
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