Thiourea can experience three intermolecular forces – London dispersion, dipole-dipole and hydrogen bonds. As with between any two molecules, Thiourea experiences London dispersion forces when attracting to another Thiourea molecule. As the electrons among the atoms’ orbits move, they tend to collect and form brief attractions to atoms of another molecule. Due to the fact that the electrons are moving, this bonding is only temporary and very weak.
The molecule also experiences dipole-dipole forces as the permanent poles of the molecule attract the pole of an opposite charge forming a weak bond which is still stronger than that of London dispersion. For example, the sulfur atom has a high negative concentration due to it unshared electron pairs so it would attract the hydrogen of another Thiourea molecule which has a positive concentration. There are actually poles located at each atom in the molecule except carbon. Unshared pairs of electrons cause the negative concentration and since hydrogen lack many electrons, they have a more positive concentration than that of other atoms. This force should not be confused with dipole-induced dipole forces which is not applicable to thiourea since this molecule can only be polar, and dipole-induced dipole occurs with polar and non-polar molecules interacting.
Lastly, thiourea can experience hydrogen bonding which is a strong bond in which the hydrogen of one molecule is attracted to the oxygen, fluorine or nitrogen of another molecule. In this case, the hydrogen would attract to the unshared pair of electrons on the nitrogen atom. This bond is a special form of the dipole-dipole forces but is much stronger than a standard dipole-dipole. It resembles an ionic bond but is weaker and can be broken by water.
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| A diagram showing the possible dipole-dipole and hydrogen bonds. |
Though the molecule lacks the two strongest bonds of ionic or metallic bonds within the molecule itself, thiourea still has several intermolecular forces and these (along with its weight) add to its high melting point. It melts (where the liquid and solid states reach equilibrium) at 182ºC and this can be compared to water which melts at 0ºC.
The polarity mentioned in previous sections is shown in the diagram above. This polarity is what results in the intermolecular forces if one thiourea molecule were to bond with another. The electron distribution in the thiourea molecule is not symmetrical horizontally and as a result, it is a polar molecule. This polarity allows for the three forces of London dispersion, dipole-dipole and hydrogen bonds to take place. As shown in the diagram, the electronegativity values of nitrogen and sulfur are greater than that of the carbon. As a result, the arrows point outward of the carbon. The hydrogen atoms have a lower electronegativity than the nitrogen atoms, hence the arrows pointing towards the nitrogen. Also represented are the negative and positive concentrations among the atoms by the stray plus and minus signs.
The differences between the electronegativity mean that each of the bonds are polar covalent and the specific differences of electronegativity between atoms are as follows:
-Between sulfur (yellow) and carbon (black): 0.1 (Very covalent)
-Between carbon and nitrogen (blue): 0.5 (Moderately covalent)
-Between nitrogen and hydrogen (white): 0.8 (Moderately covalent)
