The Chemistry of Water
Water (H2O) is the most abundant compound on Earth's surface and the medium in which essentially all biochemistry occurs. Despite its simple molecular formula — two hydrogen atoms covalently bonded to one oxygen atom — water has a collection of anomalous physical properties that make it uniquely suited to sustaining life and driving geological, atmospheric, and industrial processes.
Molecular Structure and Polarity
The oxygen atom in water is far more electronegative (3.44 on the Pauling scale) than hydrogen (2.20). This electronegativity difference makes each O-H bond strongly polar, with a partial negative charge on oxygen and a partial positive charge on each hydrogen. The molecule adopts a bent geometry (bond angle ~104.5 deg) because oxygen has two lone pairs that repel the bonding pairs according to VSEPR theory.
This bent shape is critical: it prevents the two bond dipoles from canceling, giving water a substantial permanent dipole moment of 1.85 Debye. If water were linear (like CO2), it would be nonpolar and would be a gas at room temperature — life as we know it would be impossible.
Hydrogen Bonding
Water's most consequential feature is its ability to form hydrogen bonds. The partially positive hydrogen of one water molecule is attracted to the lone pair on the oxygen of a neighboring molecule. Each water molecule can form up to four hydrogen bonds — two through its hydrogen atoms and two through its oxygen lone pairs — creating a dynamic, three-dimensional network.
Hydrogen bonds are roughly 10-40 kJ/mol, much weaker than covalent bonds (~460 kJ/mol for O-H) but much stronger than typical van der Waals forces (~1-5 kJ/mol). This intermediate strength gives water its extraordinary properties:
- High boiling point (100 degC at 1 atm) — Without hydrogen bonding, water would boil near -80 degC based on its molecular weight alone (compare H2S, which boils at -60 degC).
- High specific heat capacity (4.184 J/g-degC) — Water absorbs large amounts of heat with relatively small temperature increases. This moderates Earth's climate and stabilizes body temperature.
- High heat of vaporization (2,260 J/g) — Evaporating sweat absorbs significant heat, making sweating an efficient cooling mechanism.
- Ice floats — Upon freezing, water molecules lock into a hexagonal crystal lattice with more empty space than liquid water, giving ice a density of 0.917 g/cm3. This is anomalous; most substances are denser as solids. Floating ice insulates lakes and oceans from the top down, preventing them from freezing solid and protecting aquatic ecosystems.
Water as a Solvent
Water is called the "universal solvent" because its polarity allows it to dissolve a wider range of substances than any other common liquid. Ionic compounds (NaCl, KNO3) dissolve as water molecules surround and stabilize individual ions through ion-dipole interactions. Polar covalent molecules (ethanol, sugars) dissolve through hydrogen bonding with water.
Nonpolar molecules (oils, waxes, hydrocarbons) do not dissolve in water — they lack the polarity to compete with water's strong hydrogen-bond network. This immiscibility drives the hydrophobic effect, which is the primary force behind protein folding, cell membrane formation, and micelle assembly in soaps.
Hard Water vs. Soft Water
"Hardness" refers to the concentration of dissolved calcium (Ca2+) and magnesium (Mg2+) ions, which enter water as it percolates through limestone (CaCO3) and dolomite (CaMg(CO3)2) formations.
- Soft water: < 60 mg/L CaCO3 equivalent
- Moderately hard: 60-120 mg/L
- Hard: 120-180 mg/L
- Very hard: > 180 mg/L
Hard water causes practical problems: it forms scale (CaCO3 deposits) in pipes, boilers, and kettles; it reacts with soap to form insoluble "soap scum" (calcium stearate); and it reduces detergent efficiency. Water softeners typically use ion-exchange resins loaded with Na+ ions that swap for Ca2+ and Mg2+ ions as water passes through.
Temporary hardness (caused by dissolved calcium bicarbonate, Ca(HCO3)2) can be removed by boiling, which decomposes the bicarbonate:
Ca(HCO3)2 -> CaCO3 (precipitate) + H2O + CO2
Permanent hardness (caused by calcium sulfate or chloride) cannot be removed by boiling and requires ion exchange or reverse osmosis.
Water Purification
Municipal water treatment typically follows a multi-step process:
- Coagulation and flocculation — Aluminum sulfate (alum, Al2(SO4)3) or ferric chloride (FeCl3) is added. These salts hydrolyze in water to form gelatinous hydroxide precipitates that aggregate suspended particles into larger "flocs."
- Sedimentation — Flocs settle by gravity in large basins.
- Filtration — Water passes through sand and activated carbon beds that remove remaining particles and organic contaminants.
- Disinfection — Chlorine (Cl2 or NaOCl), chloramine (NH2Cl), ozone (O3), or UV light kills pathogens. Chlorine is the most widely used disinfectant worldwide because it provides a residual that protects water throughout the distribution system.
Reverse osmosis (RO) forces water through a semipermeable membrane with pores small enough (~0.1 nm) to block dissolved salts, producing nearly pure water. RO is essential for desalination — converting seawater (35 g/L salt) into drinking water — and for semiconductor manufacturing, where ultrapure water (18.2 megohm-cm resistivity) is required.
Water in the Environment
Water's chemical properties drive global cycles. Its high heat capacity moderates coastal climates. Its expansion upon freezing physically weathers rocks. Its solvent power carries nutrients through soil and organisms. The water cycle — evaporation, condensation, precipitation, and runoff — redistributes roughly 500,000 km3 of water annually, powered by solar energy. Understanding water chemistry is fundamental to addressing challenges from drought and contamination to climate change.