Phosphates

Structure and Hydrolysis of Phosphates

Phosphates are Mother Nature's leaving groups for nucleophilic substitution reactions. They also serve as buffers for all physiological fluids except blood (which is buffered with bicarbonate). Here are the species generally referred to as inorganic phosphate:

The pK values shown mean that when H₂PO₄^- and HPO₄^2- are present in equal concentrations in an aqueous medium, the pH will be 7.2. Actual physiological pH is about 7.4, implying a slight excess of HPO₄^2-. The pK of this latter species means that essentially no PO₄^3- exists.

The structures above are drawn differently than in typical biochemistry texts, which incorrectly show one of the P-O bonds as a double bond

In general, elements below the first full row of the Periodic Table, like phosphorus, do not form ordinary double bonds with first row elements.

Such a double bond would require interaction of a 2p orbital from the first row element, with a 3p or higher orbital from the lower element.
These orbitals have a significant size and energy mismatch, and do not overlap well.
The Lewis structure shows that this bond is really of the type we called in general chem a "dative" bond, or a "coordinate covalent" bond - the implication being that phosphorus supplies both electrons for the bond.
The result is a formal negative charge on the O and a formal positive charge on P. (Of course, the real problem is that no such thing as a two-atom, two-electron bond exists.)
To make this bond into a double bond, we would have to interact an unshared pair in a 2p orbital on oxygen with an empty 3d orbital on phosphorus, which is an energy mismatch.
Further, d orbitals lack the directional character of p orbitals. Hence, don't write PO double bonds!

The images below show the structure of inorganic phosphates, computed at the ab initio 6-31G* level of theory. The key picture is the one on the right, in which electron charge is mapped onto the surface of the molecule.

Red corresponds to high negative charge; notice that the unprotonated oxygen is significantly more negative than the others, as implied by the Lewis structure.

H₃PO₄

H₂PO₄^-

HPO₄^2-

The structures on the left show the traditional picture: the molecule drawn as composed of two-electron bonds.

For H₃PO₄ and H₂PO₄^-, the right-hand picture shows the charge distribution mapped onto the van der Waals surface; red indicates concentrations of negative charge.

The unprotonated oxygens carry the charge, just as the formal charges on the Lewis structure imply.

Finally, the right-hand structure for HPO₄^2- shows one of the molecular orbitals, in this case the HOMO - the highest occupied molecular orbital. Doesn't look anything like a "bond", does it?
Move on to the next page for a discussion of the hydrolysis of phosphates, and the so-called "energy-rich" bonds.

This page last modified 2:07 PM on Tuesday January 4th, 2011.
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