Water — structure (HOH = 104.5°), anomalous expansion, hardness (temporary - permanent), softening: Question bank

Definition Symbols and words used on this page (read first)

Nothing below is assumed — here is every piece of shorthand, in plain words:

δ+ and δ− ("delta plus / delta minus"): a partial charge, smaller than a full electron's charge. When oxygen hogs the shared electrons, its end becomes slightly negative (δ−) and each hydrogen end slightly positive (δ+). See figure s01.
μ ("mu"): the dipole moment — a single number measuring how lopsided the charge is. Bigger μ = more polar. Its unit is the D (debye); water's is μ ≈ 1.85 D.
Lone pair (LP): a pair of electrons on an atom that is not shared in a bond. Bonding pair (BP): a pair shared between two atoms to make a bond.
VSEPR (Valence Shell Electron Pair Repulsion): the rule that electron pairs around an atom spread as far apart as possible, which fixes the molecule's shape.
Electronegativity: how strongly an atom pulls shared bonding electrons toward itself; oxygen pulls harder than hydrogen, which is the origin of the δ charges.
Solubility rule used here: most sulfates and chlorides dissolve, but the carbonates of Ca²⁺ and Mg²⁺ do not.
Le Chatelier: remove a product from a reversible reaction and it keeps running forward to replace it.

Intuition The four figures on this page (read the captions, they carry the reasoning)

The cards below say things like "figure s01" or "figure s02" — here is what each picture shows, so you can map a card straight to a labelled drawing:

s01 — Water drawn bent at 104.5°: the two green O–H bond dipoles and the big coral net dipole μ ≈ 1.85 D. Use it when a card asks whether water is polar.
s02 — Oxygen with two fat butter-yellow lone-pair clouds vs two thin green bonding-pair clouds, and coral arrows showing the lone pairs squeezing the H–O–H angle down to 104.5°.
s03 — Side-by-side: the open hexagonal ice cage (empty space, low density) vs the tighter random liquid packing at 4 °C.
s04 — The density-vs-temperature curve with its peak marked at 4 °C, plus notes on why density falls on both sides of that peak.

Anchor these visually before answering — each figure below has its own caption tying it back to the cards:

s01 — Bent geometry makes the two O–H bond dipoles (green) add up into a net dipole μ ≈ 1.85 D (coral). This is why water is polar; referenced by the polarity cards.

s02 — Fat lone-pair clouds (butter) push harder than thin bonding-pair clouds (green), squeezing the angle to 104.5°. Referenced by the bond-angle cards.

s03 — Open hexagonal ice cage (left, low density) vs tighter liquid packing at 4 °C (right, high density). Referenced by the "why ice floats" cards.

s04 — Density peaks at 4 °C; it falls both when heated (thermal expansion) and when cooled below 4 °C (open cages grow). Referenced by the anomalous-expansion cards.

True or false — justify

Water's H–O–H angle is 104.5° because oxygen is sp³ hybridized and that IS the sp³ angle

False — sp³ predicts 109.5°; the extra squeeze to 104.5° comes from the two lone pairs pushing harder than bonding pairs (see figure s02), so hybridization alone does not explain the number.

If oxygen had zero lone pairs, water would still be bent

False — with 2 bonds and 0 lone pairs the electron geometry is linear, so H–O–H would open to 180°; it is the lone pairs that bend the molecule.

Ice floats because it is colder than liquid water

False — temperature is not density; ice floats because its open hexagonal cage traps empty space (figure s03), making it ~9% less dense than liquid water regardless of the exact temperature.

Water reaches maximum density exactly at its freezing point (0 °C)

False — maximum density is at 4 °C (the peak in figure s04); below 4 °C the growing open H-bonded structures make water expand again, so 0 °C water is actually less dense than 4 °C water.

A single hydrogen bond is a type of covalent bond, so it is as strong as an O–H bond

False — it is mostly an electrostatic attraction (~20 kJ/mol) with only ~10% covalent character, roughly 20× weaker than the ~460 kJ/mol O–H covalent bond.

Temporary hardness can be removed by simply adding table salt (NaCl)

False — boiling removes temporary hardness by decomposing bicarbonates to insoluble CaCO₃; NaCl adds no carbonate and does nothing to precipitate Ca²⁺.

Both temporary and permanent hardness are caused by the same Ca²⁺ and Mg²⁺ ions

True — the cations are identical; the difference is the anion they arrive with (bicarbonate → temporary; sulfate/chloride → permanent), which decides whether boiling can remove them.

Washing soda (Na₂CO₃) removes permanent hardness but boiling cannot

True — CaSO₄ and MgCl₂ stay dissolved on heating, so you must supply carbonate ions chemically to precipitate CaCO₃/MgCO₃, which Na₂CO₃ does.

Soap works better in hard water than soft water because there are more ions to react with

False — the extra Ca²⁺/Mg²⁺ waste soap as insoluble scum; soap only lathers after every hardness ion is precipitated, so hard water demands more soap for the same cleaning.

Spot the error

"Water is bent, so the two O–H bond dipoles cancel and the molecule is non-polar."

Error — in a bent shape the two O–H dipoles point partly the same way and do not cancel; they add to give a net dipole μ ≈ 1.85 D (figure s01), which is exactly why water is polar. (Contrast CO₂: linear, so its dipoles do cancel.)

"Lone pairs and bonding pairs repel equally, so NH₃, H₂O and CH₄ all keep the 109.5° angle."

Error — repulsion order is LP–LP > LP–BP > BP–BP (figure s02 shows the fatter lone-pair cloud), so each lone pair shrinks the angle: CH₄ 109.5°, NH₃ 107°, H₂O 104.5°.

"On boiling, Ca(HCO₃)₂ turns into soluble CaCO₃ and stays in the water."

Error — CaCO₃ is insoluble; per the boxed reaction

Ca (HCO_{3})_{2} Δ CaCO_{3} ↓ + H_{2} O + CO_{2} ↑

it drops out as the white scale/kettle fur, which is precisely how the Ca²⁺ leaves solution and the water softens.

"Ice sinks in most lakes, so aquatic life survives the cold at the bottom."

Error — ice floats; the floating layer insulates the liquid below and lets fish survive. If ice sank, lakes would freeze solid from the bottom up.

"The H-bond energy comes out to 24.7 kJ/mol by Coulomb's law, and that's the true value."

Error — the Coulomb estimate (~24.7 kJ/mol) is an idealized number; the real value is ≈21 kJ/mol because partial covalent character and geometry reduce the effective attraction.

"Sulfates are always insoluble, so CaSO₄ causes temporary hardness."

Error — sulfates are generally soluble (Solubility Rules); CaSO₄ stays dissolved on boiling, making it a cause of permanent (not temporary) hardness.

Why questions

Why does a lone pair repel more strongly than a bonding pair?

A lone pair is held by only one nucleus (oxygen), so its electron cloud spreads out wider and closer to O (figure s02), pushing harder on neighbouring pairs than a bonding pair shared between two nuclei.

Why is water's bent shape the cause of hydrogen bonding, not just a coincidence?

The bend leaves oxygen's lone pairs exposed on one side and the δ+ hydrogens on the other, giving water a permanent dipole; that separated charge is what lets one molecule's H attract another's lone pair.

Why does heating (above 4 °C) and cooling (below 4 °C) both decrease water's density?

Above 4 °C thermal motion expands the liquid; below 4 °C stable H-bonds build open cage structures that also expand it — 4 °C is the crossover (the peak in figure s04) where the two opposing effects balance for tightest packing.

Why does boiling remove temporary but not permanent hardness?

Boiling runs

Ca (HCO_{3})_{2} Δ CaCO_{3} ↓ + H_{2} O + CO_{2} ↑

so the carbonate precipitates Ca²⁺; but sulfates/chlorides have no such heat-triggered breakdown, so those ions remain dissolved.

Why does adding heat shift the bicarbonate equilibrium toward precipitation?

By Le Chatelier's Principle, driving off gaseous CO₂ and forming solid CaCO₃ both remove products from solution, pulling the reaction forward and stripping Ca²⁺ out of the water.

Why does the 1:1 stoichiometry of Ca(HCO₃)₂ → CaCO₃ matter for softening calculations?

One mole of dissolved bicarbonate salt yields exactly one mole of solid carbonate, so moles removed equal moles present — no ratio bookkeeping needed to find the precipitate.

Why is water called an "anomalous" liquid instead of just "unusual"?

"Anomalous" flags that it breaks the general rule (solids denser than their liquids); its density peak at 4 °C and expansion on freezing are genuine exceptions, driven by H-bond geometry rather than normal thermal behaviour.

Edge cases

If a water sample has bicarbonates AND sulfates of calcium, is it temporary or permanent hard?

Both — boiling removes the bicarbonate (temporary) portion, but the sulfate portion (permanent) survives, so full softening still needs a chemical like Na₂CO₃ or an ion-exchange resin.

What happens to water hardness if the sample contains only NaCl (no Ca²⁺/Mg²⁺)?

It is not hard — Na⁺ does not form scum with soap; hardness is defined by Ca²⁺/Mg²⁺ (and sometimes Fe²⁺), not by dissolved salts in general.

At exactly 0 °C, can water be both liquid and solid, and which is denser?

Yes — at 0 °C ice (0.917 g/cm³) and liquid water (~0.998 g/cm³) coexist, and the liquid is denser, so the ice floats on it.

Does distilled (ion-free) water lather with soap?

Yes, readily — with no Ca²⁺/Mg²⁺ to precipitate the soap, essentially all of it goes to forming cleaning micelles, so it is effectively "soft" water.

If you keep boiling permanently hard water for hours, does it eventually soften?

No — no amount of boiling decomposes CaSO₄ or MgCl₂; the ions stay dissolved, so the water remains hard until you treat it chemically.

Would water still show maximum density at 4 °C if hydrogen bonding did not exist?

No — without H-bonds there would be no open cage tendency on cooling, so water would simply grow denser all the way to freezing like a normal liquid, with no 4 °C peak and no floating ice.