"Chemical Bonding and Molecular Structure" — Practice

Question 1

Predict the shape and hybridisation of SF$_4$ and explain where the lone pair sits.

Accepted Answer

Sulphur has six valence electrons, four used in bonds to fluorine, leaving one lone pair. The steric number is 4 plus 1, that is 5, so the electron geometry is trigonal bipyramidal and the hybridisation is $sp^{3}d$. The lone pair takes an equatorial position, where it faces only two neighbours at $90^\circ$ rather than three. Ignoring it leaves a see-saw shape, with the axial F to S to F angle bent slightly below $180^\circ$.

Question 2

Why is XeF$_4$ square planar while SF$_4$ is see-saw, when both have lone pairs on the central atom?

Accepted Answer

Because their steric numbers differ. Xenon in XeF$_4$ has eight valence electrons, four in bonds and four as two lone pairs, giving a steric number of 6 and an octahedral electron geometry with $sp^{3}d^{2}$ hybridisation. The two lone pairs take opposite axial positions to be as far apart as possible, leaving four fluorines in a plane. SF$_4$ has steric number 5, a trigonal bipyramidal geometry and a single equatorial lone pair, which gives the see-saw.

Question 3

Arrange BF$_3$, NH$_3$ and H$_2$O by bond angle, from largest to smallest, and explain.

Accepted Answer

Boron trifluoride is $120^\circ$, ammonia $107^\circ$ and water 104.5. Boron trifluoride is trigonal planar with steric number 3, $sp^{2}$ hybridised and with no lone pairs, so it keeps the ideal angle. Ammonia and water are both $sp^{3}$ with an ideal $109.5^\circ$, compressed by one and two lone pairs respectively, since lone pair repulsion exceeds bond pair repulsion. Comparing across different steric numbers requires stating both the ideal geometry and the compression.

Question 4

Calculate the bond order of N$_2$, N$_2^{+}$ and N$_2^{-}$ and predict which has the shortest bond.

Accepted Answer

Nitrogen has 14 electrons, giving 10 bonding and 4 antibonding, so the bond order is 3. N$_2^{+}$ has 13 electrons, removing one from the bonding $\sigma 2p$, giving a bond order of 2.5. N$_2^{-}$ has 15, adding one to an antibonding pi orbital, also giving 2.5. Neutral N$_2$ has the highest bond order and therefore the shortest and strongest bond. Note that for nitrogen the electron removed is bonding, unlike oxygen where it is antibonding.

Question 5

Explain why He$_2$ does not exist but He$_2^{+}$ does.

Accepted Answer

Helium has two electrons, so He$_2$ has four: two fill the bonding $\sigma 1s$ and two fill the antibonding $\sigma^{*}1s$. The bond order is $(2-2)/2$, which is zero, so there is no net bonding and the molecule dissociates. He$_2^{+}$ has three electrons, two bonding and one antibonding, giving a bond order of 0.5. That is weak but non-zero, and He$_2^{+}$ has indeed been detected spectroscopically.

Question 6

Why is O$_2$ paramagnetic when its Lewis structure shows all electrons paired?

Accepted Answer

Because the Lewis model has no way to represent degenerate orbitals. Oxygen's 16 electrons fill up to the two degenerate $\pi^{*}$ orbitals, which receive one electron each with parallel spins by Hund's rule rather than pairing in one of them. Two unpaired electrons remain, so oxygen is paramagnetic, and liquid oxygen is visibly attracted to a magnet. This prediction is the strongest evidence for molecular orbital theory over the Lewis and valence bond pictures.

Question 7

Both CO$_2$ and SO$_2$ contain polar bonds, yet only one is a polar molecule. Explain.

Accepted Answer

Carbon dioxide is linear, with steric number 2 on carbon and no lone pairs, so its two bond dipoles are equal and opposite and cancel exactly, giving a zero net dipole moment. Sulphur dioxide has a lone pair on sulphur, steric number 3 and a bent shape, so its bond dipoles do not cancel and add to a resultant of about 1.6 D. Shape decides polarity, not bond polarity, which is why symmetry must always be checked before answering.

Question 8

Explain why the dipole moment of NH$_3$ is 1.47 D while that of NF$_3$ is only 0.24 D.

Accepted Answer

Both are trigonal pyramidal with one lone pair, so shape is not the difference. In ammonia, nitrogen is the more electronegative atom, so each bond dipole points from hydrogen towards nitrogen, in the same direction as the lone pair dipole, and the four contributions reinforce. In nitrogen trifluoride, fluorine is more electronegative than nitrogen, so the bond dipoles point away from nitrogen and oppose the lone pair dipole, leaving only a small resultant.

Question 9

Which is more covalent, AgCl or NaCl, and why?

Accepted Answer

Silver chloride. Both cations carry a single positive charge and their radii are comparable, but Ag$^{+}$ has a pseudo noble gas configuration of $4d^{10}$ rather than the true noble gas core of Na$^{+}$. A $d^{10}$ shell shields the nucleus poorly, so the effective nuclear charge at the surface of Ag$^{+}$ is higher and its polarising power is much greater. It distorts the chloride cloud strongly, giving substantial covalent character, which is why silver chloride is insoluble in water while sodium chloride is freely soluble.

Question 10

Why does magnesium oxide have a far higher melting point than sodium chloride?

Accepted Answer

Because lattice enthalpy scales with the product of the ionic charges and inversely with the internuclear distance. Magnesium oxide is built from doubly charged Mg$^{2+}$ and O$_2^{-}$ ions, so the charge product is four times that of the singly charged ions in sodium chloride, and both ions are also smaller. Its lattice enthalpy is near 3800 kJ mol$^{-1}$ against sodium chloride's 788, and its melting point exceeds 2800 °C against 801.

Question 11

Ortho-nitrophenol boils lower than para-nitrophenol. Explain.

Accepted Answer

In the ortho isomer the nitro group sits adjacent to the hydroxyl group, close enough for the hydroxyl hydrogen to hydrogen bond to a nitro oxygen within the same molecule. That intramolecular bond, sometimes called chelation, uses up both the donor and the acceptor, so few intermolecular hydrogen bonds form and the molecules separate easily. In the para isomer the groups are too far apart for an internal bond, so extensive intermolecular hydrogen bonding raises the boiling point.

Question 12

Explain why ice is less dense than liquid water, and why this matters.

Accepted Answer

In ice, each water molecule hydrogen bonds to four others in a fixed tetrahedral arrangement, which forces an open cage structure containing substantial empty space. On melting, that lattice partly collapses and molecules pack closer, so liquid water is denser and reaches its maximum density near 4 °C. Ice therefore floats, which insulates the water beneath and allows aquatic life to survive winters. Almost every other substance is denser as a solid than as a liquid.

Practice — "Chemical Bonding and Molecular Structure"

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