3.2.7 · D5p-Block
Question bank — Group 16 (Oxygen family) — allotropes of O (O₂, O₃); ozone chemistry, ozone layer
0. Symbols and pictures you need first
Before the traps, here is a small toolbox so no symbol below is a mystery. Read this once, then the reveals will make sense from line one.
Figure 1 — the MO energy ladder that makes O₂ paramagnetic. Read the ladder bottom (low energy, bonding) to top (high energy, antibonding). Count the arrows: the topmost two electrons land singly in the two boxes — that is the unpaired pair every trap below refers to.
Figure 2 — the ozone resonance hybrid. Left and middle: the two Lewis drawings (one bond single, one double). Right: the real molecule — a bent shape with the π electron density smeared equally over both bonds, so both are identical (bond order 1.5).
Figure 3 — why is uphill. The reaction-coordinate curve climbs to a higher product energy: . Ozone sits on a shelf above O₂, storing that energy — which is why it later rolls back down, releasing .
Figure 4 — the CFC chlorine chain. Follow the loop: Cl attacks O₃, becomes ClO·, then hands its O to a stray atom and is reborn as Cl. The arrow returning to the start is why one atom wrecks thousands of O₃.
Figure 5 — why the layer shields us. UV-absorption curves: O₂ only catches the shortest, most energetic UV; O₃ additionally mops up the UV-B/UV-C band that would otherwise reach the ground and damage DNA.
True or false — justify
Ozone and dioxygen are different compounds.
False. They are allotropes — different molecular forms of the same element oxygen, not different compounds (a compound needs two or more different elements). See the parent note.
Because O₂ is drawn as O=O, it must be diamagnetic.
False. The Lewis picture cannot show antibonding orbitals; Molecular Orbital Theory places two unpaired electrons in the orbitals (see Figure 1), so liquid O₂ is paramagnetic (it clings to a magnet).
Ozone's two O–O bonds have different lengths, one being a double bond.
False. The single/double drawing is only one resonance form; the real molecule is the resonance hybrid (Figure 2, right) with both bonds equal, each bond order 1.5 and length ≈128 pm.
Ozone is more thermodynamically stable than O₂ because it has more atoms.
False. Ozone's formation is endothermic (, Figure 3), so it stores extra energy and is less stable — that instability is exactly why it decomposes and oxidises so readily.
A single chlorine atom from a CFC destroys just one ozone molecule.
False. Cl acts as a catalyst: it is regenerated in (the return arrow in Figure 4), so one Cl atom cycles through and destroys thousands of O₃ molecules — see Free Radical Chain Reactions.
The ozone layer works by physically blocking UV like a solid shield.
False. It works by absorbing UV chemically (Figure 5): the same photon-driven cycle that splits O₂ and O₃ converts UV energy into heat, so harmful radiation is consumed, not reflected.
Ground-level ozone protects us from the Sun, just like stratospheric ozone.
False. Only stratospheric ozone shields us; tropospheric (ground-level) ozone is a toxic oxidising pollutant in photochemical smog — see Environmental Chemistry — Air Pollution.
The central oxygen atom in ozone is sp hybridised because ozone is linear.
False. Ozone is bent (~117°): the central O carries 2 bond pairs + 1 lone pair, giving roughly sp² hybridisation, not sp, and not linear (Figure 2).
O₃ has a higher bond order than O₂.
False. O₂ has bond order 2; each O–O in O₃ is only 1.5. Lower bond order means the O₃ bond is weaker and longer than the O₂ bond.
Nascent oxygen and molecular O₂ are equally good oxidisers.
False. Nascent is a single, highly reactive atomic species that oxidises far more aggressively than the stable, paired-up O₂ molecule — that is why the released drives ozone's oxidising power (Oxidation and Reducing Agents).
Spot the error
"Ozone is prepared by heating O₂ strongly until it converts to O₃."
Wrong method. Ozone is made by a silent electric discharge through O₂; strong heat would just decompose any ozone formed (the reaction is endothermic, so a controlled discharge, not thermal heating, drives it).
"Bond order = ."
Missing the ÷2. With = bonding electrons and = antibonding electrons, the correct MOT formula is ; each pair of net bonding electrons contributes one bond, so you divide by two.
"In the ozone resonance , the molecule flips rapidly between the two structures."
No flipping happens. Resonance structures are not real states the molecule switches between; the true molecule is a single fixed hybrid (Figure 2, right) whose electron density is a blend of both drawings.
" releases energy, so ozone is stable."
Sign error. This reaction is endothermic (, Figure 3); it absorbs energy, making ozone thermodynamically unstable.
"CFCs destroy ozone by reacting with it directly as whole molecules."
They don't. CFCs are inert until UV cleaves a C–Cl bond releasing a Cl radical; it is the radical, not the intact CFC, that attacks ozone (Figure 4).
"For the KI test, oxidises iodide to iodate."
Wrong product. Ozone oxidises iodide to iodine (I₂): ; the liberated I₂ is then titrated with thiosulphate.
"O₂ is paramagnetic because it has an odd number of electrons."
Wrong cause. O₂ has 16 electrons (even); its paramagnetism comes from two unpaired electrons sitting singly in the two degenerate orbitals (Hund's rule, Figure 1), not from an odd count.
Why questions
Why does ozone act as a powerful oxidising agent while O₂ is comparatively mild?
Ozone easily decomposes as , and the released nascent oxygen is a hungry atomic oxidiser; O₂ has no such easily-released reactive atom.
Why is the O–O bond in ozone (128 pm) between that of O₂ (121 pm) and hydrogen peroxide's O–O single bond (148 pm)?
Because its bond order (1.5) is between a double bond (2) and a single bond (1); intermediate bond strength gives intermediate length.
Why does one Cl atom cause disproportionate ozone loss compared to its tiny amount?
Cl is a catalyst in a chain (Figure 4): it converts O₃ to O₂ and is then regenerated, so the same atom repeats the destruction cycle thousands of times before it is finally removed.
Why can't a simple Lewis structure explain O₂'s magnetism, but MOT can?
Lewis structures only pair electrons in bonds and lone pairs; they have no concept of antibonding orbitals, whereas MOT shows the two unpaired electrons occupying antibonding (Figure 1).
Why is the natural stratospheric Chapman cycle described as a dynamic equilibrium rather than a one-way reaction?
Ozone is being continuously formed () and destroyed () at balanced rates, so its concentration stays roughly steady while UV is being absorbed throughout.
Why does the "tailing of mercury" happen when ozone is present?
Ozone oxidises the mercury surface, forming a film that makes mercury stick to the glass and lose its clean meniscus (a qualitative test for O₃).
Edge cases
If all the antibonding electrons of O₂ were removed to make , would the bond be stronger or weaker?
Stronger. Removing antibonding () electrons raises the bond order (fewer destabilising electrons in ), so the bond becomes shorter and stronger than in neutral O₂.
Is there any temperature/pressure at which O₂ and O₃ stop being "allotropes"?
No — allotropy is defined by their different bonding in the same physical state; regardless of conditions they remain two molecular forms of the element oxygen.
What happens to the Chapman balance if incoming UV suddenly vanished?
With no UV, O₂ would no longer be split into atoms, so no new O₃ would form; the layer's steady state collapses because its formation step depends entirely on UV.
Could ground-level ozone ever be beneficial as a germicide?
Yes in controlled use (water/air disinfection) its oxidising power kills microbes; but as an uncontrolled component of smog in the air we breathe it is a harmful pollutant — context decides.
If a resonance form of ozone showed both bonds identical already, would we still need resonance?
We need resonance precisely because no single Lewis structure can show two identical 1.5-order bonds; the hybrid of the two unequal drawings is what produces the equal, delocalised bonds.
Recall One-line takeaways
- Allotropes = same element, different molecular form ::: not different compounds.
- O₂ paramagnetic ::: 2 unpaired e⁻ in (MOT, not Lewis — Figure 1).
- O₃ bonds ::: equal, bond order 1.5, resonance hybrid, ~117° bent (Figure 2).
- Ozone reactive ::: releases nascent ; endothermic formation ⇒ unstable (Figure 3).
- Cl in CFC chain ::: catalyst, regenerated, destroys thousands of O₃ (Figure 4).
- Stratospheric O₃ shields ::: tropospheric O₃ pollutes (Figure 5).