← Unit 2: How do chemical reactions shape the natural world?
How do chemicals interact with water?
the explanation of the properties of water (including high boiling point, high specific heat capacity, surface tension and the density of ice relative to liquid water) and the role of water as a solvent for polar and ionic substances, including the use of solubility rules to predict precipitation reactions and write ionic equations
A focused VCE Chemistry Unit 2 answer on the chemistry of water. Covers hydrogen bonding and how it explains water's anomalous physical properties, how water dissolves ionic and polar molecular substances, the use of solubility rules to predict precipitation, and writing balanced ionic and net ionic equations.
Have a quick question? Jump to the Q&A page
What this dot point is asking
VCAA wants you to explain the properties of water using its molecular structure and hydrogen bonding, to describe how water acts as a solvent for ionic and polar molecular substances, and to use solubility rules to predict the products of precipitation reactions and write full, ionic and net ionic equations.
The answer
Why water has unusual properties
Water is a small, bent, polar molecule with two O-H bonds and two lone pairs on the oxygen. Each H2O molecule can hydrogen bond to up to four neighbours (donating two H atoms and accepting two via its lone pairs). The resulting hydrogen-bond network produces several anomalous properties:
| Property | What you see | Bonding explanation |
|---|---|---|
| High boiling point (100 deg C) | Liquid at room temperature, despite low molar mass | Hydrogen bonds must be broken in addition to weaker IMFs |
| High specific heat capacity (4.18 J g^-1 K^-1) | Water heats and cools slowly | Energy goes into breaking H-bonds, not just into kinetic energy |
| High enthalpy of vaporisation | Sweating is an effective coolant | Many H-bonds break per gram during evaporation |
| Surface tension | Water beads on a leaf; insects walk on water | H-bonds at the surface pull inward, contracting the surface |
| Density of ice less than liquid water | Ice floats | Open hexagonal hydrogen-bonded lattice in ice has more empty space than liquid water |
| Universal solvent for polar/ionic | Dissolves salts, sugars, alcohols | Polar O-H bonds let water orient around ions and dipoles |
Water as a solvent
Ionic compounds. When an ionic solid (e.g. NaCl) is added to water, the dipoles of water orient around the ions. The partial negative O attracts cations (Na^+), the partial positive H atoms attract anions (Cl^-). Each ion is surrounded by a hydration shell of water molecules. The ions separate from the lattice (dissociation) and become aquated:
NaCl(s) -> Na^+(aq) + Cl^-(aq)
If the energy released by hydration (hydration enthalpy) is comparable to or greater than the lattice energy holding the ions together, the salt dissolves. If not, the salt is insoluble.
Polar molecular substances. Substances such as sugar, ethanol or ammonia dissolve in water because they can hydrogen bond (or at least form dipole-dipole interactions) with water. They do not dissociate into ions; they dissolve as intact molecules.
Non-polar substances (oil, hexane, iodine in water) do not dissolve because the water-water hydrogen-bond network is not compensated by any equally favourable interaction with the solute. They form a separate layer.
The shorthand: "like dissolves like". Polar/ionic dissolves in polar solvents; non-polar dissolves in non-polar solvents.
Solubility rules
A workable VCE-level set of solubility rules for ionic compounds in water:
| Always (mostly) soluble | Insoluble exceptions |
|---|---|
| Group 1 cation salts (Li^+, Na^+, K^+, etc.) | None worth memorising |
| Ammonium (NH4^+) salts | None |
| Nitrates (NO3^-) | None |
| Acetates / ethanoates | None |
| Chlorides, bromides, iodides | Ag^+, Pb^2+, Hg2^2+ are insoluble |
| Sulfates (SO4^2-) | Ag^+, Pb^2+, Ba^2+, Ca^2+ (slightly), Sr^2+ insoluble |
| Mostly insoluble | Soluble exceptions |
|---|---|
| Carbonates (CO3^2-) | Group 1 cations, NH4^+ |
| Phosphates (PO4^3-) | Group 1 cations, NH4^+ |
| Hydroxides (OH^-) | Group 1 cations, NH4^+, Ca^2+ (slightly), Ba^2+, Sr^2+ |
| Sulfides (S^2-) | Group 1 cations, NH4^+ |
Use the VCAA data book version on the day. The rules above are sufficient for most VCE questions.
Predicting precipitation and writing equations
When two solutions are mixed, swap the partners and check each potential product against the solubility rules. If either product is insoluble, a precipitate forms.
Procedure for writing a net ionic equation:
- Write the full molecular equation with states (aq, s, l, g).
- Split all (aq) ionic compounds into their separate ions; leave (s), (l), (g) and weak acids/bases intact.
- Cancel spectator ions that appear unchanged on both sides.
- Check that mass and charge balance.
Worked example
A student mixes 50 mL of 0.10 mol L^-1 lead(II) nitrate solution with 50 mL of 0.10 mol L^-1 potassium iodide solution. Predict the precipitate and write the net ionic equation.
Potential products by swapping partners: PbI2 and KNO3. Solubility rules: KNO3 is soluble (Group 1 + nitrate); PbI2 is insoluble (lead iodide is one of the famous exceptions). So PbI2 is the precipitate (a bright yellow solid).
Full: Pb(NO3)2(aq) + 2KI(aq) -> PbI2(s) + 2KNO3(aq)
Ionic: Pb^2+(aq) + 2NO3^-(aq) + 2K^+(aq) + 2I^-(aq) -> PbI2(s) + 2K^+(aq) + 2NO3^-(aq)
Spectators: K^+ and NO3^- cancel.
Net ionic: Pb^2+(aq) + 2I^-(aq) -> PbI2(s)
Common traps
Writing H2O as the reactant for dissolving. Water surrounds the ions but is not consumed. Use (aq) on the ions and (s) on the solid; do not put H2O in the equation unless it is genuinely a reactant or product.
Treating weak acids and weak bases as fully ionic in an ionic equation. Weak acids (CH3COOH, HF, etc.) and weak bases (NH3) stay molecular. Only strong electrolytes split.
Forgetting states. A net ionic equation without (aq), (s), (l), (g) is incomplete.
Cancelling ions that have changed in number. A K^+ on both sides at 2 each cancels; a Pb^2+ that appears as 1 on the left and is now inside the PbI2 precipitate is not a spectator.
Listing PbI2 as soluble because most iodides are. Read the exceptions: Ag, Pb and Hg(I) halides are insoluble. Always check both sides of the rule.
Saying ice is less dense because it is colder. Most substances are denser when solid. Ice is the standout exception because its open hydrogen-bonded lattice traps empty space.
In one sentence
Water's small bent polar shape and hydrogen-bonding give it high boiling point, high specific heat, high surface tension and a less-dense solid form, the same hydrogen bonding lets it dissolve polar and ionic substances, and applying the solubility rules to a mixture of two solutions predicts whether a precipitate will form and what the net ionic equation looks like.
Past exam questions, worked
Real questions from past VCAA papers on this dot point, with our answer explainer.
2024 VCE4 marksExplain in terms of bonding why (a) ice is less dense than liquid water, and (b) water has an unusually high specific heat capacity. Why are these properties important for aquatic life?Show worked answer →
A 4-mark answer needs the molecular explanation and the biological link.
(a) Ice less dense than water: in liquid water, hydrogen bonds form and break rapidly and molecules pack closely. In ice, each water molecule forms 4 hydrogen bonds to 4 neighbours in a fixed, open tetrahedral lattice. The lattice has large empty spaces, so the same number of molecules occupies a larger volume. Density (mass/volume) is therefore lower.
(b) High specific heat capacity: heating water raises kinetic energy, but a significant fraction of the added energy goes into breaking hydrogen bonds (potential energy) rather than into translational/rotational kinetic energy. More energy per gram is required for each 1 degree rise.
Aquatic life: ice floats on top of lakes and insulates the liquid water below, allowing fish and aquatic plants to survive winter. The high specific heat capacity stabilises ocean and lake temperatures, buffering organisms from rapid temperature swings.
2025 VCE4 marksWhen aqueous solutions of silver nitrate and sodium chloride are mixed, a white precipitate forms. (a) Write the full balanced equation, (b) the ionic equation, and (c) the net ionic equation. (d) Use the solubility rules to explain which product is the precipitate.Show worked answer →
A 4-mark answer needs all three forms of the equation and the rules cited.
(a) Full equation:
AgNO3(aq) + NaCl(aq) -> AgCl(s) + NaNO3(aq)
(b) Ionic equation (separate all aqueous species into their ions):
Ag^+(aq) + NO3^-(aq) + Na^+(aq) + Cl^-(aq) -> AgCl(s) + Na^+(aq) + NO3^-(aq)
(c) Net ionic equation (cancel spectator ions Na^+ and NO3^-):
Ag^+(aq) + Cl^-(aq) -> AgCl(s)
(d) Solubility rules state that most nitrates are soluble (so NaNO3 stays in solution) and most chlorides are soluble except those of silver, lead and mercury(I). AgCl is therefore the insoluble white precipitate. The precipitation is driven by the strong ionic attraction in the AgCl lattice exceeding the hydration energies of Ag^+ and Cl^- on their own.
Related dot points
- expressing the concentration of solutions (mol L^-1, g L^-1, %m/v, %m/m, %v/v and ppm) including dilution calculations, and the Brønsted-Lowry model of acids and bases including conjugate acid-base pairs, the distinction between strong and weak (and concentrated and dilute) acids and bases, and the calculation of pH from [H+]
A focused VCE Chemistry Unit 2 answer on concentration and acid-base chemistry. Covers concentration units (mol L^-1, g L^-1, %m/v, %m/m, %v/v, ppm) and dilution calculations, the Brønsted-Lowry model with conjugate acid-base pairs, strong vs weak and concentrated vs dilute, and the calculation of pH from [H+].
- the principles of volumetric analysis including acid-base and redox titrations, the use of primary and secondary standard solutions and indicators, and stoichiometric calculations including back-titration to determine the concentration or amount of analyte
A focused VCE Chemistry Unit 2 answer on volumetric analysis. Covers acid-base and redox titrations, primary and secondary standards, the choice of indicator from titration curves, the c1V1 / c2V2 / mole-ratio workflow, and back-titration for samples that react slowly or with excess.