Shapes of molecules and intermolecular forces


Leaving Certificate Chemistry (Shapes of molecules and intermolecular forces) Note on Shapes of molecules and intermolecular forces, created by eimearkelly3 on 06/08/2013.
Note by eimearkelly3, updated more than 1 year ago
Created by eimearkelly3 almost 11 years ago

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Molecules are formed when atoms are joined together by covalent bondsThe simplest molecule is diatomic (two molecules - linear)

SHAPESLinearV-shapedTrigonal PlanarPyramidalTetrahedral

Electron repulsion theoryThe electron pairs in the valence shell of a central atom repel each other and end up as far apart as geometrically possible.Lone pairs have a greater repelling ability than bonding pairs.Since lone pairs are closer to the nucleus of the central atom, they are closer to each other, and so their mutual repulsion is greater than the mutual repulsion between bonding pairs.lone pair : lone pair > lone pair : bonding pair > bonding pair : bonding pairThe result of this is that where lone pairs and bonding pairs occur in the valence shell of the central atom, distortion in the expected arrangement is found, and the bond angles differ to what might otherwise be expected.

Symmetry and polarityA polar molecule musta) contain polar covalent bondsb) not be symmetricalCENTRES OF + AND - CHARGESometimes a molecule can be non-polar overall even though the individual bonds are polar. This occurs when the molecule has a high degree of symmetry with a geometric centre.e.g. beryllium hydride --> the beryllium atom is the centre of positive charge, since beryllium is less electronegative than hydrogen. Being midway between the two hydrogen atoms, it is also the centre of negative charge. Thus both centres of charge coincide and the molecule is non-polar overall.SEPERATED CENTRES OF + AND - CHARGEMolecules of water for example, contain polar covalent bonds  but do not contain the same degree of polarity as say beryllium hydride. Consequently the centres of positive and negative charge do not coincide (permanent dipoles exist), and the molecule is polar overall.

Intramolecular bonding -- within the atom

Intermolecular forces -- between molecules Van der Waal's forces Dipole-Dipole interactions Hydrogen bonding These forces are weaker than covalent bonds but affect properties such as boiling point (molecules are more difficult to seperate and a greater amount of heat needs to be provided to allow molecules to escape the main bulk of the liquid). The stronger the attractive forces, the higher the boiling point of the liquid.

VAN DER WAALS -- Johannes Van Der WaalsThese forces are much weaker than covalent bonds and can themselves vary in strength.They are weak attractive forces caused by the movement of electrons within a molecule. Electrons move randomly within a bond so that at a particular point in time they may be nearer to one atom than the other. This creates a temporary polarity (temporary dipole) in the molecule.If two molecules with similar temporary dipoles happen to be oriented with opposite charges directed at each other, an attractive force will exist between the two molecules.Sometimes, a temporary dipole in one molecule will induce a similar dipole in a neighbouring, with an effect similar to the above. The combined effect of these temporary and induced dipoles is to produce weak attractive forces between neighbouring molecules, resulting in increased boiling points.The grerater the number of electrons in a molecule, the greater the number of possible temporary dipoles. Consequently, the extent of the intermolecular attraction is greater. This explains why oxygen gas (with 16 electrons per molecule) has a higher boiling point than hydrogen (with only 2 electrons per molecule).

Dipole-Dipole interactions between polar moleculesSimilar to Van Der Waal's forces in that the negative end of one dipole is attracted to the positive end of another, but they differ in that the dipoles are permanent due to the polarity in the molecule.Because they are permanent, they are stronger than Van der Waal's forces.e.g. ethene (C2H4), and methanol (HCHO) would be expected to have boiling points owing to the similar size of their relative molecular massses.This would be the case if there were only Van der Waals forces between their respective molecules. However, ethene boils at  169K while methanol boils at 252K. The difference arises because of methanol's stronger intermoleclar forces.

HYDROGEN BONDINGstronger than other types of dipole-dipole interactionsHydrogen bonding is a special type of dipole-dipole interaction, which occurs when hydrogen is bonded to small, highly electronegative atoms such ass O, N or F.-In waterIn water molecules, the O-H bond is highly polar owing to the large electronegativity difference between hydrogen and oxygen. OXYGEN carries a partial NEGATIVE charge and HYDROGEN carries a partial POSITIVE charge.These strong intermolecular forces cause a much higher boiling point in water, 373K. Hydrogen bonding is also the main intermolecular force that holds moleccules of water together in ice crystals.

Water and hydrogen sulfideHydrogen sulfide may be expected to have a higher boiling point than water owing to its greater molecular mass. However it boils at 211 KBecause:The H-S bond is much less polar than the O-H bond in water.The partial negative charge is more diffuse and thus less effective on the large sulfur atom than on the smaller oxygen atom.Thus there is no hydrogen bondinng between the H2S molecules, and the dipole-dipole interactions are much weaker. The large difference in boiling point shows the strength of the h-bonding in water.

Shapes of molecules

Intramolecular bonding and intermolecular forces


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