3.1.7 · D5Hydrogen and s-Block
Question bank — Alkali metals (Group 1) — physical - chemical properties, anomaly of Li, diagonal Li-Mg
Before we start, one word we will lean on: reducing agent means "a substance that donates electrons to something else" — it hands over electrons and gets oxidised itself. Keep that picture (an electron being pushed away from the metal) in mind; half the traps below hinge on it.
True or false — justify
Down Group 1, atomic radius increases but first ionization enthalpy decreases — these two trends are just coincidences that happen to point opposite ways.
False — they are the same cause. A bigger radius means the lone outer electron sits farther from the nucleus, so it is held more weakly, so less energy is needed to remove it. One trend drives the other.
Because Li has the highest ionization enthalpy in Group 1, it must be the weakest reducing agent in water.
False — in water we must add hydration enthalpy. Tiny Li⁺ releases so much energy on being surrounded by water that it overcompensates its high I.E., making Li the strongest reducing agent in aqueous solution ( V).
Metallic bonding gets stronger as you go down Group 1 because the atoms get heavier.
False — bond strength depends on how tightly the delocalised electron sea grips the cations, not on mass. Bigger atoms spread the same single valence electron over a larger volume, so the grip weakens and melting points fall down the group.
Density increases smoothly from Li to Cs.
False — the trend is Li < Na < K < Rb < Cs. Potassium is the anomaly: adding the 4th shell inflates its volume faster than its mass grows, so its density dips below sodium's before recovering.
All alkali metals burn in oxygen to give the same type of oxide.
False — Li gives the normal oxide , Na gives the peroxide , and K, Rb, Cs give superoxides (). Larger cations stabilise larger, more loosely-held anions.
Lithium is the most reactive alkali metal toward water because it has the smallest atom.
False — reactivity with water follows Li < Na < K < Rb < Cs. Cs loses its electron most easily (lowest I.E.), so it reacts most violently; Li is the calmest of the group.
The diagonal similarity between Li and Mg means they are in the same group.
False — Li is Group 1, Mg is Group 2. They resemble each other because charge/size ratio (polarising power) is similar across the diagonal, not because of group membership. See 3.1.09-Diagonal-relationship-general.
Alkali metal halides are all freely soluble in water.
Almost — LiF is the exception. Li⁺ is so small and F⁻ so small that their lattice energy is huge, and this outweighs hydration, making LiF only sparingly soluble.
Spot the error
"Since Z (nuclear charge) increases down Group 1, the outer electron is pulled in harder, so ionization energy rises."
The error is ignoring distance and shielding. barely rises because each new full inner shell shields strongly; meanwhile the electron sits a whole shell farther out. Distance wins, so I.E. falls.
" Li⁺ is bigger than Li because it has lost an electron and so has fewer repulsions pushing electrons outward."
Backwards. Losing the lone electron removes an entire shell, so Li⁺ is much smaller than Li. Fewer repulsions and an unshielded nucleus pull the remaining electrons in tighter.
"K is less dense than Na, therefore K is a lighter atom than Na."
Confuses density with atomic mass. K is heavier (39 vs 23) but its volume balloons even more, so mass/volume drops. Density is a ratio, not a mass.
"Cs forms a superoxide because Cs is the most reactive, so it grabs the most oxygen."
Reactivity is not the cause. The reason is lattice stability: the large Cs⁺ cation packs stably with the large superoxide anion, whereas small Li⁺ needs the small oxide.
"Flame colours differ because heavier metals simply glow brighter."
The colour comes from the energy gap an excited electron falls back across: . Different metals have different gaps, hence different wavelengths (colours) — brightness is unrelated.
"Alkali metals are stored in kerosene because they rust in air like iron does."
Not rusting — they react instantly and violently with moisture and oxygen. Kerosene keeps water and air away entirely; "rust" (slow oxidation) is a different, gentler process.
Why questions
Why is Li often called "anomalous" within its own group?
Its atom and ion are exceptionally small, giving it very high polarising power. This makes Li behave more covalently and resemble Mg (diagonal) rather than its heavier Group-1 siblings — see 4.2.03-Ionic-vs-covalent-character.
Why does the same factor (small ionic size) make Li both the strongest reducing agent in water and the least reactive toward water?
Different arenas. In water, small Li⁺'s enormous hydration enthalpy drives the overall (thermodynamics of the finished ion). The rate of the violent surface reaction is governed by how fast the metal sheds electrons and the heat that ignites H₂ — Li's high I.E. and high melting point make that sluggish.
Why must we use a full thermodynamic cycle (sublimation + I.E. + hydration) to rank reducing power in solution, instead of I.E. alone?
Because describes the metal going all the way from solid to hydrated ion. I.E. is only the middle gas-phase step; ignoring sublimation and hydration gives the wrong ranking, which is exactly why Li surprises us. See 5.3.02-Standard-reduction-potentials.
Why do all these radius, I.E., melting-point, and reactivity trends "line up" so neatly for Group 1?
Because Group 1 varies just one thing cleanly — a single electron getting farther out and better shielded down the group. With everything else held constant, the periodic trends all trace back to that one moving electron.
Why do alkali metals reduce so strongly compared to Group 2 metals?
They need to release only one loosely held electron to reach a noble-gas core, and that lone electron is weakly bound. Group 2 must remove two electrons and holds them more tightly (higher ).
Edge cases
Fr (francium) sits below Cs — do the trends simply continue, and can we test them easily?
The trends should continue (largest radius, lowest I.E., most reactive), but Fr is intensely radioactive and short-lived, so it exists only in trace, fleeting amounts — the predictions are extrapolations, not lab-measured.
What is the "degenerate" case: an alkali cation like Na⁺ — does it still want to donate an electron?
No. Once it is Na⁺ it has the stable noble-gas core it wanted; there is no lonely outer electron left to give. Reducing power belongs to the neutral metal, not its ion.
If you compared a hypothetical alkali metal with zero shielding, would I.E. still fall down the group?
No. Shielding is precisely what lets stay nearly flat while distance grows. Without it, rising would tighten the grip and I.E. could rise — the falling trend depends on the shielding that real inner shells provide.
At the boundary between "peroxide" and "superoxide" formers (Na vs K), what single property flips the product?
Cation size. Na⁺ is still small enough to favour the peroxide lattice; K⁺ crosses the size threshold where the larger, singly-charged superoxide gives the more stable lattice.
Recall Rapid self-quiz
Down Group 1, I.E. goes ::: down (bigger radius, weak grip on the lone electron). Strongest reducing agent in water ::: Li, thanks to its huge hydration enthalpy. Density anomaly ::: K is less dense than Na. Oxide types Li / Na / K ::: normal oxide / peroxide / superoxide. Li resembles which diagonal neighbour ::: Mg.