Unit 8 worksheet 6 Lewis acids & bases
1. Boron trifluoride, BF3, can react with potassium fluoride to form the salt potassium tetrafluoroborate,
KBF4.
Explain why this can be considered an acid-base reaction and deduce the shapes of both boron
trifluoride and the tetrafluoroborate ion and the F―B―F bond angles in each case.
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2. Aluminium chloride, AlCl3, can act as a halogen carrier in organic chemistry, for example, benzene
reacts with chlorine in the presence of aluminium chloride to form chlorobenzene and hydrogen
chloride. It works by breaking the Cl―Cl bond to form AlCl4–and a positive chloride ion, Cl+, which can
then react with benzene. Explain why the reaction between aluminium chloride and chlorine can be
considered to be an acid-base reaction.
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3. Explain why all ligands and all nucleophiles can also be considered as Lewis bases.
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4. The cyanide ion, CN− is a good ligand. It is poisonous as it reacts irreversibly with transition metal ions
in the body.
(a) Draw the Lewis structure of the cyanide ion.
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(b) Use the concept of formal charge to suggest why it is the carbon atom of the cyanide ion that bonds
to the transition metal ion and not the nitrogen atom.
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(c) Using information from the spectrochemical series suggest why drinking an aqueous solution of
iron(II) sulfate is an effective antidote to cyanide poisoning provided it is taken immediately after the
cyanide has been ingested.
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5. Aluminium chloride. AlCl3 readily dimerises to form Al2Cl6. Explain this dimerisation reaction in terms
of Lewis acid and base theory clearly identifying the Lewis acid and the Lewis base.
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Answers
1. BF3 is accepting a pair of electrons and thus acting as a Lewis acid. F– is donating the pair of electrons
and thus is acting as a Lewis base. (It is worth noting that BF3 is an ‘electron deficient compound’ as
there are only three pairs of electrons around the central boron atom. It can readily accept one more
pair of electrons to complete the octet.)
BF3 : trigonal planar, 120o. BF4– : regular tetrahedron, 109.5o.
2. AlCl3 only contains three electron pairs around the central aluminium atom. When it acts as a halogen
carrier it completes its octet by accepting a pair of electrons from one of the chlorine atoms in the
chlorine molecule causing the Cl―Cl bond to break and form AlCl4― and Cl+. It is therefore acting as a
Lewis acid and the chlorine molecule is acting as a Lewis base.
3. All ligands and all nucleophiles possess at least one pair of non-bonding electrons and they react by
donating this pair of electrons and are thus acting as Lewis bases. Ligands donate them to transition
metal ions and nucleophiles donate them to an electron rich carbon atom.
4. (a)
(b) Both the carbon and the nitrogen atoms in the cyanide ion possess a non-bonding pair of electrons.
The formal charge on the carbon atom = 4 − 2 − (1/2 x 6) = −1; the formal charge on the nitrogen atom =
5 − 2 − (1/2 x 6) = 0. Since the carbon atom has a more negative formal charge it is more likely to donate
its non-bonding pair of electrons to the transition metal ion.
(c) From the spectrochemical series it can be seen that cyanide ion ligands are more electron dense than
water ligands as they causes greater splitting of the d orbitals. Because the cyanide ion is more electron
dense it is a stronger ligand than water. If aqueous iron(II) ions, [Fe(H2O)6]2+, are added to cyanide ions
the cyanide ions replace the six water ligands irreversibly to form [Fe(CN)6]4−. This means the cyanide are
no longer free to form coordinate bonds with other transition metal ions in the body.
5. The aluminium atom in AlCl3 is electron deficient as it only has six electrons in its outer shell. It can
therefore act as a Lewis acid accepting a non-bonding pair of electrons from one of the chlorine atoms
from another molecule of aluminium chloride. At the same time one of its own chlorine molecules also
donates a pair of electrons to the other aluminium atom. AlCl3 is therefore acting as both a Lewis acid
and a Lewis base