Published by Patrick Mutisya · 14 days ago
Each water molecule consists of one oxygen atom covalently bonded to two hydrogen atoms. The molecule adopts a bent shape with an angle of about 104.5° between the H–O–H bonds. Oxygen is more electronegative than hydrogen, so the O–H bonds are polar: the oxygen end carries a partial negative charge (δ‑) and the hydrogen ends carry partial positive charges (δ+).
When water molecules approach one another, the δ+ hydrogen of one molecule is attracted to the δ‑ oxygen of a neighbouring molecule. This electrostatic attraction is called a hydrogen bond. A single water molecule can form up to four hydrogen bonds – two as a donor (through its two hydrogens) and two as an acceptor (through the two lone pairs on oxygen).
\$\text{H}2\text{O} \; \cdots \; \text{H}2\text{O}\$
Hydrogen bonds are relatively weak compared with covalent bonds (≈ 20 kJ mol⁻¹), but because each molecule can participate in several of them, a large number of bonds exist throughout liquid water, giving rise to its distinctive properties.
Water’s polarity allows it to dissolve a wide range of ionic and polar substances. The process can be described by the dissociation of an ionic compound into its constituent ions, which become surrounded by water molecules (hydration). For example:
\$\text{NaCl}{(s)} \rightarrow \text{Na}^+{(aq)} + \text{Cl}^-_{(aq)}\$
In living organisms, this property enables water to transport nutrients, waste products, gases, and signalling molecules throughout cells and tissues.
The specific heat capacity (c) of water is 4.18 J g⁻¹ °C⁻¹, one of the highest among common substances. The energy required to raise the temperature of 1 g of water by 1 °C is therefore:
\$q = m \times c \times \Delta T = 1\;\text{g} \times 4.18\;\text{J g}^{-1}\!\!^\circ\text{C}^{-1} \times 1^\circ\text{C} = 4.18\;\text{J}\$
Because a large amount of heat must be absorbed or released before water’s temperature changes appreciably, water buffers temperature fluctuations in organisms and environments, maintaining relatively stable internal conditions (homeostasis).
Water requires 2260 J g⁻¹ to change from liquid to vapour at 100 °C (its latent heat of vaporisation). This high value reflects the energy needed to break the extensive hydrogen‑bond network.
Biologically, the large latent heat is exploited in evaporative cooling mechanisms such as sweating and panting, where the conversion of liquid water to vapour removes excess heat from the body.
| Property | Value (SI units) | Biological Relevance |
|---|---|---|
| Polarity | Dipole moment ≈ 1.85 D | Enables dissolution of ions and polar molecules; essential for metabolic reactions. |
| Specific heat capacity (c) | 4.18 J g⁻¹ °C⁻¹ | Buffers temperature changes in cells, blood, and aquatic habitats. |
| Latent heat of vaporisation (Lᵥ) | 2260 J g⁻¹ at 100 °C | Provides efficient cooling via evaporation (sweat, respiration). |
| Hydrogen‑bonding capacity | Up to 4 H‑bonds per molecule | Creates cohesive network that influences viscosity, surface tension, and thermal properties. |