Pictures of the Day CH320M328M

A good way to predict relative acid strength is to look at the conjugate base anions produced upon deprotonation of a set of acids, and predict which anion will be the most stable anion. The more stable the conjugate base anion, the more acidic the parent acid. The best way to stabilize an anion is to spread the negative charge around. Shown above on the left is the ethoxide anion (CH3CH2O-), with the negative charge (darkest red color) localized entirely on the oxygen atom. The molecule in the middle is the acetate anion (CH3CO2-), and notice how the negative charge (lighter red color) is distributed over the two oxygen atoms, as explained by the two relevant resonance structures you should be able to draw for this species. Spreading the charge due to this so-called "resonance effect" is extremely stabilizing compared to the ethoxide anion so it should be no surprise that the parent acid of the acetate anion, namely acetic acid, has a pKa of 4.75 and is significantly more acidic than the parent acid of ethoxide anion, namely ethanol, having a pKa of around 16. The molecule on the right is the trifluoroacetate anion (CF3CO2-), and here, like the acetate anion in the middle, there is a resonance effect helping to distribute the negative charge over the two oxygen atoms. However, with the trifluoroacetate anion, there are also three electronegative fluorine atoms attached nearby, and these pull some electron density through the sigma bonding network away from the oxygen atoms, thereby spreading out the negative charge (lightest red color) even further. This effect, called the "inductive effect" is stabilizing to anions such as the trifouoroacetate anion so the parent acid, trifluoroacetic acid is the strongest acid of the three by far, having a pKa of -0.18.

Shown on the left is the BF3 molecule. With three bonds to fluorine (sp2 hybridization), and no lone pairs, there remains one 2p-orbital that is not hybridized and empty. Thus the boron atom needs a lone pair of electrons to give it an octet of electrons. Also, fluorine is highly electronegative, withdrawing electron density from the boron atom. This is represented in the electrostatic potential model at the upper-left, with flourine atoms in yellow-orange (partial negative charge) and boron in blue (partial positive charge). For all these reasons, the molecule BF3 is a good acceptor of electrons and therefore a good Lewis acid. Ammonia (NH3), shown in the middle, has a lone pair of electrons, and since nitrogen is more electronegative than hydrogen, the nitrogen atom has a partial negative charge (red color). In this diagram, NH3 acts as a Lewis base,when it donates its lone pair of electrons to BF3. BF3 acts as a Lewis acid when it accepts the lone pair of electrons that NH3 donates. This reaction fills BF3's empty 2p-orbital, and now boron is sp3 hybridized when previously (as BF3) it was sp2 hybridized. Note the negative charge on the boron half of the molecule (red color) and positive charge on the ammonia half of the molecule (blue color).