Natural Waters: Composition and Properties

Introduction

No natural water is pure H₂O. Rainwater carries dissolved gases and traces of sea salt and dust; rivers pick up ions as they flow over rocks; the ocean has accumulated salts over billions of years. The composition of a water decides whether it scales kettles, supports trout or tastes salty, and it sets the stage for all the aquatic chemistry in the rest of this unit. This page surveys what is dissolved in natural waters, where it comes from, and why water itself is such an unusual solvent.

Core explanation

Major ions. Seven ions account for more than 95% of dissolved solids in most natural waters: the cations Na⁺, K⁺, Ca²⁺ and Mg²⁺ and the anions Cl⁻, SO₄²⁻ and HCO₃⁻. Their proportions differ sharply between rivers and the ocean:

Ion Average river water / mg L⁻¹ Seawater / mg L⁻¹ --- --- --- HCO₃⁻ about 58 about 140 Ca²⁺ about 15 about 410 Mg²⁺ about 4 about 1290 Na⁺ about 6 about 10 800 Cl⁻ about 8 about 19 400 SO₄²⁻ about 11 about 2700

River water is dominated by Ca²⁺ and HCO₃⁻ from weathering of limestone and silicates: CaCO₃ + CO₂ + H₂O → Ca²⁺ + 2HCO₃⁻. Seawater is dominated by Na⁺ and Cl⁻. Although rivers supply more calcium than sodium, marine organisms remove calcium and bicarbonate as CaCO₃ shells, whereas Na⁺ and Cl⁻ are almost unreactive and accumulate. Their long residence times (tens of millions of years for Na⁺) explain seawater's saltiness, around 35 g of salts per kilogram. The ratios of major ions in the open ocean are nearly constant (Marcet's principle), so salinity can be found from conductivity alone.

Dissolved gases. Gas solubility follows Henry's law, c = K H × p, where K H depends on the gas and temperature. O₂ is sparingly soluble: fresh water at 20 °C saturated with air holds only about 9 mg L⁻¹ of O₂ (about 0.28 mmol L⁻¹), falling to about 7.5 mg L⁻¹ at 30 °C. Warm water holds less gas because dissolution is exothermic. CO₂ is about thirty times more soluble than O₂ because it reacts partly to form carbonic acid, as explored in the next page on the carbonate system.

Hardness. Hardness is caused mainly by Ca²⁺ and Mg²⁺. It is expressed as the equivalent mass of CaCO₃: hardness (mg L⁻¹ as CaCO₃) = 2.50 × [Ca²⁺ in mg L⁻¹] + 4.12 × [Mg²⁺ in mg L⁻¹]. Water below about 60 mg L⁻¹ as CaCO₃ is classed as soft, above about 180 as very hard. Temporary hardness, associated with bicarbonate, is removed on boiling as CaCO₃ precipitates (limescale): Ca²⁺ + 2HCO₃⁻ → CaCO₃ + CO₂ + H₂O. Permanent hardness, associated with sulfate and chloride, is not.

Unusual properties of water. Extensive hydrogen bonding gives water an exceptionally high heat capacity (4.18 J g⁻¹ K⁻¹), a high enthalpy of vaporisation, a high surface tension and a boiling point far above those of H₂S or H₂Se. Its high relative permittivity (about 78 at 25 °C) weakens attraction between ions, making it an excellent solvent for salts. Liquid water reaches maximum density at about 4 °C, and ice is less dense than liquid, so lakes freeze from the top and deep water stays near 4 °C in winter.

Formulae

Henry's law: c = K H × p. Hardness (as CaCO₃) = 2.50 [Ca²⁺] + 4.12 [Mg²⁺], all in mg L⁻¹. The factors are 100.1/40.08 and 100.1/24.31.

Step-by-step reasoning

To interpret a water analysis:

1. Check the charge balance: sum of cation equivalents should roughly equal sum of anion equivalents. 2. Identify the dominant cation and anion to classify the water (for example Ca–HCO₃ or Na–Cl type). 3. Relate the type to its source: carbonate rock, evaporites, sea spray or pollution. 4. Calculate hardness from Ca²⁺ and Mg²⁺. 5. Compare dissolved oxygen with the saturation value at the measured temperature.

Visual explanation

Draw two pie charts of cation equivalents: river water dominated by a large Ca²⁺ slice, seawater dominated by Na⁺. Beneath, sketch a density-versus-temperature curve for water peaking at 4 °C, and a lake cross-section in winter showing ice at the surface over 4 °C water at the bottom.

Real-world analogy

The ocean behaves like a bank account where deposits of every ion arrive by river, but only some are withdrawn. Calcium and bicarbonate are spent quickly on shells, while sodium and chloride are hardly ever withdrawn. Over time the account fills with the ions that nobody spends.

Real-world example

Water from chalk aquifers in south-east England often has hardness above 250 mg L⁻¹ as CaCO₃, causing limescale in kettles and boilers, whereas upland reservoir water in Scotland and Wales, draining granite and peat, is typically below 50 mg L⁻¹ and can be slightly acidic and corrosive to pipes.

Why?

Why does warm water hold less oxygen than cold water? Dissolving a gas is exothermic and decreases entropy, as molecules move from the freely moving gas into an ordered solvation shell. Raising the temperature therefore shifts the equilibrium back towards the gas phase, reducing the Henry's law constant.

Common misconception

"Seawater is salty because rivers carry mostly sodium chloride." River water actually contains more calcium and bicarbonate. Seawater is sodium chloride rich because those ions are rarely removed, whereas calcium and bicarbonate are extracted by marine organisms and precipitated as carbonate sediments.

Worked example

Question: A water sample contains 88 mg L⁻¹ Ca²⁺ and 12 mg L⁻¹ Mg²⁺. Calculate its hardness as CaCO₃ and classify it.

Reasoning: Hardness = 2.50 × 88 + 4.12 × 12 = 220 + 49 = 269 mg L⁻¹ as CaCO₃. This exceeds 180 mg L⁻¹.

Answer: About 270 mg L⁻¹ as CaCO₃; very hard water.

Quick check

1. Why can a lake's deep water stay liquid near 4 °C when the surface freezes over in winter? Answer: Water is densest near 4 °C, so this water sinks, while less dense colder water and ice float on top.

Exam focus

Learn the seven major ions, the reason for the river–sea contrast, Henry's law, and the hardness calculation with its conversion factors. Link water's anomalous properties explicitly to hydrogen bonding and explain temporary versus permanent hardness with an equation.

Advanced insight

Henry's law constants must be corrected for salinity: gases are less soluble in seawater than in fresh water, the "salting-out" effect, because ions bind water molecules in hydration shells and leave less free solvent. Seawater at 20 °C holds about 7.4 mg L⁻¹ of O₂, compared with about 9.1 mg L⁻¹ in fresh water.

Summary

Natural waters are dominated by seven major ions. Rivers are rich in Ca²⁺ and HCO₃⁻ from rock weathering, while seawater is rich in Na⁺ and Cl⁻, which accumulate over long residence times. Gas solubility follows Henry's law and falls with temperature. Hardness from Ca²⁺ and Mg²⁺ is expressed as CaCO₃ and may be temporary or permanent. Hydrogen bonding gives water its high heat capacity, solvent power and density maximum at 4 °C.

Practice questions

1. Explain why calcium and bicarbonate dominate river water. Answer: Rain containing dissolved CO₂ weathers limestone and silicate rocks, releasing Ca²⁺ and HCO₃⁻: CaCO₃ + CO₂ + H₂O → Ca²⁺ + 2HCO₃⁻. 2. Write an equation to show how temporary hardness is removed by boiling. Answer: Ca²⁺ + 2HCO₃⁻ → CaCO₃ + CO₂ + H₂O; the carbonate precipitates as limescale. 3. The Henry's law constant for O₂ at 25 °C is 1.3 × 10⁻³ mol L⁻¹ atm⁻¹. Find the dissolved O₂ concentration in water in contact with air (p(O₂) = 0.21 atm). Answer: c = 1.3 × 10⁻³ × 0.21 = 2.7 × 10⁻⁴ mol L⁻¹, about 8.7 mg L⁻¹. 4. Give two properties of water that result from hydrogen bonding and their environmental importance. Answer: High heat capacity moderates climate and water temperatures; the density maximum at 4 °C means lakes freeze from the top, protecting aquatic life.