The more I look critically at the IB Chemistry Diploma programme the more I find that is either simply wrong or at best only a simplified half-truth. This is not to criticise the IB programme as such as all other pre-university programmes suffer from the same problems. How aware are teachers about the truth or otherwise of some of the ‘accepted’ chemistry on the programme?

Clearly sometimes we have to give a simplified version as the mathematics is too difficult for students at this level  - an example of this is describing ethyne as containing two pi bonds at right angles to each other. Other times the models used contain so many flaws that they are wrong even though both students and teachers should be very much aware of the errors. For example, the way in which ionic bonding is often taught. Other areas are less clear cut. Hydrogen bonding falls into one of these less clear cut areas.

Students are taught that hydrogen bonding occurs when hydrogen is directly bonded to one of the three most electronegative atoms, fluorine, oxygen and nitrogen. In a sense fluorine is not so important as there is only one possible compound, hydrogen fluoride, HF. Hydrogen bonded directly to oxygen and nitrogen provides many important examples of hydrogen bonding such as the anomalous behaviour of water (left) and the ‘glue’ holding the double helices together in DNA. However a critical look at this ‘definition’ causes a problem straightaway. According to Table 7 in the IB Chemistry Data Booklet chlorine (3.2) is more electronegative than nitrogen (3.0) and bromine (3.0) has the same value so hydrogen bonded to chlorine or bromine should also show hydrogen bonding. Yet in the classic graph of the boiling points of the hydrides for groups 4,5,6 and 7 (below) only the abnormally high boiling points of water, ammonia and hydrogen fluoride are highlighted as being due to hydrogen bonding.

In fact hydrogen bonding is just a more extreme case of dipole-dipole interactions between molecules. A normal van der Waals’ type of attraction between non-polar molecules is in the order of 1 kJ mol^-1, dipole-dipole interactions are a little more than this and hydrogen bonding can vary between about 10 - 25 kJ mol^-1.  IUPAC have this year (2011) changed their definition of what constitutes a hydrogen bond thus rendering our text books somewhat obsolete. This change, which was published in Pure and Applied Chemistry, 2011, Vol 83, No. 8, p 1637-1641, bases the definition on theoretical and experimental evidence obtained during the past century. The full text can be read as a pdf file.  The new definition is “The hydrogen bond is an attractive interaction between a hydrogen atom from a molecule or a molecular fragment X–H in which X is more electronegative than H, and an atom or a group of atoms in the same or a different molecule, in which there is evidence of bond formation”. This clearly still includes hydrogen bonded directly to fluorine, oxygen and nitrogen but now it also includes, for example, hydrogen bonded directly to carbon. This bond at IB level is usually considered to be non-polar but carbon (2.6) is more electronegative than hydrogen (2.2) so the bond is in fact polar. For an intermolecular force to be classified as a hydrogen bond it has to meet certain experimental conditions which are listed in detail in the article. These  include, bond angles, infrared absorptions and NMR shifts.

Probably this is another example where books and pre-university syllabi will stick with the simplified version, i.e. with the ‘hydrogen bonded directly to fluorine, oxygen and nitrogen’ definition, and yet teachers should be aware that it is nowhere near this simple in reality.