Water
Water is an inorganic compound with the chemical formula H₂O, comprising two hydrogen atoms covalently bonded to one oxygen atom in a bent molecular structure.[1][2] It appears as a colorless, odorless, and tasteless liquid at standard temperature and pressure, with a density of approximately 1 g/cm³ at 4°C, a melting point of 0°C, and a boiling point of 100°C at 1 atm.[1] Water exhibits anomalous properties, including maximum density at 4°C, expansion upon freezing, high specific heat capacity, and elevated heat of vaporization, which arise from extensive hydrogen bonding between molecules.[1] These characteristics enable water to serve as the universal solvent for polar and ionic substances, facilitating chemical reactions and transport in biological systems.[3] On Earth, water constitutes about 71% of the planet's surface, predominantly in oceans as saline solution comprising roughly 97% of the total hydrosphere, with the remainder as freshwater in glaciers, groundwater, and surface bodies.[4] Essential for all known life forms, water forms the medium for metabolic processes, regulates temperature via its thermal properties, and supports structural integrity in cells through cohesion and adhesion.[5] Its polarity and hydrogen bonding underpin its role in dissolving nutrients, enabling enzymatic activity, and maintaining homeostasis in organisms.[6]
Etymology and Nomenclature
Linguistic and Historical Origins
The English term "water" originates from Old English wæter, attested in texts from the 9th century CE, referring to the liquid essential for life and its wetting properties. This form evolved from Proto-West Germanic watar and Proto-Germanic watōr, a root shared across Germanic languages, including Old Saxon watar, Old Norse vatn, Dutch water, and German Wasser.[7][8] Linguistically, the Proto-Germanic term descends from Proto-Indo-European *wódr̥ (or variant *wédōr), an ablaut form implying "water" or "wet," reconstructed through comparative analysis of cognates in other Indo-European branches, such as Slavic voda (e.g., Russian voda), Baltic vanduo (Lithuanian), and Tocharian wär. This root reflects a semantic focus on moisture and flowing liquids, distinct from alternative PIE terms like *h₂ep- for bodies of water or *h₂ékʷeh₂ yielding Latin aqua and thus Romance equivalents like French eau and Spanish agua.[9][10] Historically, the term's significance emerged in the early 20th century when Czech linguist Bedřich Hrozný deciphered Hittite cuneiform in 1915, identifying wa-a-tar (from PIE *wódr̥) in texts dating to circa 1600–1200 BCE from ancient Anatolia, providing pivotal evidence for the Indo-European language family's extent beyond Europe. This cognate underscored water's conceptual primacy in proto-languages, as a basic environmental and ritual element, though non-Indo-European ancient terms like Sumerian a (circa 3000 BCE) or Egyptian mw (from Pyramid Texts, circa 2400 BCE) represent unrelated, independently developed designations for the substance.[11][12]Chemical Composition
Molecular Structure and Bonding
The water molecule, with the chemical formula H₂O, consists of a single oxygen atom covalently bonded to two hydrogen atoms.[13] These bonds are polar covalent, arising from the electronegativity difference between oxygen (3.44 on the Pauling scale) and hydrogen (2.20), which results in a partial negative charge on the oxygen atom and partial positive charges on the hydrogen atoms.[13] The oxygen atom features two lone pairs of electrons, leading to sp³ hybridization and a tetrahedral electron-pair geometry according to valence shell electron pair repulsion (VSEPR) theory.[14] The molecular geometry of water is bent or V-shaped, with an H–O–H bond angle of 104.5°.[14] [15] This angle is smaller than the ideal tetrahedral value of 109.5° due to greater repulsion between the lone pairs than between bonding pairs.[13] The bent structure and polar bonds produce a net dipole moment of approximately 1.85 Debye, making water a polar molecule.[16] Beyond intramolecular covalent bonding, water exhibits intermolecular hydrogen bonding, where the partially positive hydrogen atom of one molecule forms an electrostatic attraction with the partially negative oxygen atom of a neighboring molecule.[16] Each water molecule can participate in up to four hydrogen bonds—two as a donor via its hydrogens and two as an acceptor via its lone pairs—forming a dynamic, fluctuating tetrahedral network in the liquid state.[17] This hydrogen bonding network, with an average of about 3.5 bonds per molecule, underlies many of water's anomalous properties, such as its high boiling point and cohesion.[18] The bond energy of a single hydrogen bond in water is approximately 20 kJ/mol, significantly weaker than the intramolecular O–H covalent bond at 460 kJ/mol.[19]Isotopes and Variants
Water primarily consists of molecules formed from the stable isotopes of hydrogen and oxygen. Hydrogen has two stable isotopes: protium (^1H), comprising approximately 99.98% of hydrogen atoms in natural waters, and deuterium (^2H or D), at about 0.02%.[20] Oxygen has three stable isotopes: ^16O (99.63%), ^17O (0.0375%), and ^18O (0.1995%). The most abundant form is thus ^1H_2^16O, but natural waters contain trace amounts of isotopic variants due to these abundances. Tritium (^3H or T), a radioactive hydrogen isotope with a half-life of 12.32 years, occurs in negligible quantities from cosmic rays and nuclear processes. Deuterated forms represent key variants. Heavy water, or deuterium oxide (D_2O), features two deuterium atoms, resulting in a molecular mass of 20.0276 g/mol, higher density of 1.107 g/mL at 20°C, a melting point of 3.82°C, and boiling point of 101.4°C—elevated compared to ordinary water due to stronger hydrogen bonding from the heavier isotope's reduced zero-point energy.[21] [22] Semiheavy water (HDO) contains one protium and one deuterium atom; in typical water, HDO molecules constitute about 1 in 3,200, far outnumbering pure D_2O, and exhibit intermediate properties blending those of H_2O and D_2O.[23] Oxygen isotopic substitution, such as in H_2^18O, increases density slightly (e.g., by ~0.1% for 1% ^18O enrichment) but has minimal impact on phase transitions relative to hydrogen variants.[24] These variants arise from primordial nucleosynthesis for deuterium and oxygen isotopes, with distributions maintained by fractionation processes like evaporation and condensation, which preferentially enrich lighter isotopes in vapor phases.[24] Heavy water is industrially produced via electrolysis or distillation exploiting the ~8°C boiling point difference, achieving >99.8% purity.[25] It serves as a neutron moderator in pressurized heavy-water reactors (e.g., CANDU designs), slowing neutrons without significant absorption to sustain fission in natural uranium fuel, unlike light water which requires enrichment. Tritiated water (e.g., HTO or T_2O) incorporates tritium, behaving chemically like ordinary water but with added radioactivity from beta emission (average energy 5.7 keV). Pure T_2O has a density of ~1.21 g/mL, melting point of 4.48°C, and boiling point of 101.51°C, though environmental forms are dilute mixtures.[26] Its biological half-life in humans is ~10 days, distributing uniformly in body water, with health risks assessed via dose limits (e.g., 7,000 Bq/L drinking water standard in some regulations).[27] Tritium production occurs in reactors via neutron capture on deuterium or lithium.[28] Isotopic studies of water variants aid hydrology, paleoclimatology, and forensics by tracking fractionation signatures.[23]Physical Properties
States of Matter and Phase Transitions
Water exists in solid, liquid, and gaseous states depending on temperature and pressure. At standard atmospheric pressure of 101.325 kPa, liquid water freezes into ice at 0 °C (273.15 K) and boils into vapor at 100 °C (373.15 K).[29] The solid phase, ice Ih, exhibits a lower density than liquid water, with ice at 0 °C having a density of approximately 916.7 kg/m³ compared to 999.8 kg/m³ for liquid water at the same temperature, causing ice to float on water.[30] This density inversion arises from the open hexagonal lattice structure in ice, stabilized by hydrogen bonding, which expands upon solidification.[31] Phase transitions between these states involve latent heat absorption or release without temperature change. The heat of fusion for melting ice is 333.55 J/g, while the heat of vaporization at 100 °C is 2257 J/g.[32] Sublimation, the direct transition from solid to gas, occurs below the triple point, as seen in dry ice but applicable to water under low pressure. The triple point, where solid, liquid, and gas coexist in equilibrium, occurs at 0.01 °C (273.16 K) and 611.657 Pa.[33] Above the critical point at 374 °C (647 K) and 22.064 MPa (218 atm), water enters a supercritical state indistinguishable as liquid or gas, exhibiting properties of both.[34] Liquid water displays a density maximum at approximately 4 °C (277 K), decreasing upon further cooling due to enhanced hydrogen bonding that increases molecular volume, contributing to the density anomaly.[30] Under elevated pressures, water forms multiple polymorphs of ice, such as ice II, III, and others, with at least 18 distinct phases identified, each stable in specific pressure-temperature regimes revealed by the water phase diagram.[35] These transitions underscore water's unique thermodynamic behavior, driven by its polar molecular structure and hydrogen bonding network.Thermodynamic Characteristics
Water possesses several distinctive thermodynamic properties arising primarily from intermolecular hydrogen bonding, which imparts higher energy requirements for changes in molecular arrangement compared to non-hydrogen-bonded liquids of similar molecular weight. The isobaric specific heat capacity (Cp) of liquid water at 25°C and standard atmospheric pressure is 4.184 J/(g·K), or 75.3 J/(mol·K), allowing it to store substantial thermal energy with minimal temperature rise; this value exceeds that of ethanol (2.44 J/(g·K)) and acetone (2.15 J/(g·K)) by factors of approximately 1.7 and 1.9, respectively.[36][37] The difference stems from the cooperative disruption of hydrogen bonds during vibrational excitation, as confirmed by molecular dynamics simulations linking bond network entropy to heat capacity anomalies.[38] Latent heats of phase transitions are notably elevated: the enthalpy of fusion (ΔfusH) for ice at 0°C is 333.55 J/g (6.01 kJ/mol), while the enthalpy of vaporization (ΔvapH) at 100°C is 2256.4 kJ/kg (40.65 kJ/mol), values roughly double those expected for non-associated liquids like methane derivatives.[39] These high latent heats reflect the energy needed to overcome tetrahedral hydrogen-bonded structures in the liquid and solid phases, enabling water to moderate environmental temperatures effectively, as observed in oceanic heat retention.[40] The triple point, where solid, liquid, and vapor phases coexist in equilibrium, occurs at 0.01°C (273.16 K) and 611.657 Pa, marking the boundary beyond which ice sublimes directly under reduced pressure.[41] Water's critical point, at which liquid and vapor phases become indistinguishable, is reached at 374°C (647.1 K) and 22.064 MPa (218.3 atm), higher than for comparable non-polar fluids due to persistent hydrogen bonding suppressing supercritical mixing until extreme conditions.[39] Thermally, liquid water exhibits negative expansion below 4°C, with maximum density of 999.975 kg/m³ at 3.98°C, an anomaly driven by the collapse of open-cage hydrogen-bond networks into denser configurations upon cooling, contrasting the typical contraction of liquids.[41] The coefficient of isobaric thermal expansion (αp) averages 2.57 × 10−4 K−1 near 20°C, while isothermal compressibility (κT) is low at 4.59 × 10−10 Pa−1, indicating resistance to volume change under pressure.[41] Thermal conductivity peaks at 0.68 W/(m·K) around 130°C, facilitating efficient heat transfer in natural systems.[42]| Property | Value at Standard Conditions | Notes |
|---|---|---|
| Specific heat capacity (Cp, liquid, 25°C) | 4.184 J/(g·K) | High due to hydrogen bond disruption[36] |
| Enthalpy of fusion (0°C) | 333.55 J/g | Energy to break ice lattice bonds[39] |
| Enthalpy of vaporization (100°C) | 2256 kJ/kg | Overcomes full network in liquid[39] |
| Thermal expansion coefficient (αp, 20°C) | 2.57 × 10−4 K−1 | Anomalous below 4°C[41] |
| Isothermal compressibility (κT, 25°C) | 4.59 × 10−10 Pa−1 | Low, enhancing incompressibility[41] |
Mechanical and Optical Properties
Water's mechanical properties stem from its hydrogen-bonded network, conferring behaviors distinct from many other liquids. The dynamic viscosity of liquid water at 25°C is 0.89 mPa·s, reflecting resistance to shear flow that diminishes with temperature as thermal energy overcomes intermolecular attractions.[44] Surface tension at the water-air interface stands at 72.0 mN/m at 25°C, a value elevated by cohesive forces that minimize surface area, manifesting in capillary rise and droplet sphericity.[45] Compressibility is low, quantified by a bulk modulus of 2.2 GPa under ambient conditions, signifying that pressures on the order of hundreds of megapascals are required for measurable volume reduction, unlike gases.[46] Optically, pure water transmits visible light (400–700 nm) with high transparency, absorbing less than 0.01% per meter in the blue-green range, which enables deep penetration in clear aquatic environments.[47] Its refractive index for visible wavelengths is 1.333 at 20°C, varying slightly with wavelength (higher for shorter wavelengths) and causing phenomena such as mirages and the bending of light at interfaces.[48] Absorption intensifies in the ultraviolet (strong below 200 nm due to electronic excitations) and infrared (peaking at vibrational modes around 3 μm and beyond), rendering water opaque in those spectra despite visible clarity.[49]Chemical Properties
Reactivity and Ionization
Water exhibits limited chemical reactivity under ambient conditions, attributable to the high bond dissociation energy of its O-H bonds, approximately 498 kJ/mol for the first bond.[50] This stability contrasts with its ability to participate in specific reactions, particularly with electropositive elements and certain oxides. For instance, alkali metals react exothermically with water, displacing hydrogen gas and forming hydroxides: $2\mathrm{Na} + 2\mathrm{H_2O} \rightarrow 2\mathrm{NaOH} + \mathrm{H_2}, with reaction vigor increasing from lithium to cesium due to decreasing ionization energies and lattice energies of the metals.[51] Alkaline earth metals, such as magnesium and calcium, react more slowly, often requiring heating or steam for complete reaction, as in \mathrm{Ca} + 2\mathrm{H_2O} \rightarrow \mathrm{Ca(OH)_2} + \mathrm{H_2}.[52] Water also engages in hydrolysis reactions with non-metal oxides and halides, demonstrating its role as a nucleophile. Carbon dioxide reacts partially to form carbonic acid: \mathrm{CO_2 + H_2O \rightleftharpoons H_2CO_3}, influencing ocean acidity.[53] Similarly, phosphorus pentachloride hydrolyzes: \mathrm{PCl_5 + 4H_2O \rightarrow H_3PO_4 + 5HCl}, releasing HCl gas. These reactions underscore water's amphoteric character, allowing it to act as both a Lewis base (donating electron pairs to electrophiles) and, in specific contexts, facilitating proton transfer.[54] In terms of ionization, water undergoes autoionization: $2\mathrm{H_2O \rightleftharpoons H_3O^+ + OH^-}, governed by the ion product constant K_w = [\mathrm{H_3O^+}][\mathrm{OH^-}] = 1.0 \times 10^{-14} at 25 °C.[55] [53] This equilibrium yields equal concentrations of hydronium and hydroxide ions in pure water, approximately $1.0 \times 10^{-7} M each, corresponding to a neutral pH of 7.00. The pK_a for water acting as an acid (\mathrm{H_2O \rightleftharpoons H^+ + OH^-}) is approximately 15.7, reflecting the low extent of deprotonation due to the strong O-H bond and solvation effects; this value derives from K_a = K_w / [\mathrm{H_2O}], where [\mathrm{H_2O}] \approx 55.5 M.[56] [57] The gas-phase ionization energy of the water molecule, required to remove an electron from the highest occupied molecular orbital, measures 12.62 ± 0.01 eV, as determined by photoelectron spectroscopy.[58] In liquid water, solvation lowers the vertical ionization energy to about 11.67 eV, facilitating processes like radiolysis but still indicating high energy barriers under thermal conditions.[59] These properties collectively position water as a poor conductor in pure form, with ionization primarily driven by external fields or impurities rather than intrinsic thermal dissociation.Electrical Conductivity and Electrolysis
Pure water exhibits very low electrical conductivity due to its limited autoionization, which produces hydronium (H₃O⁺) and hydroxide (OH⁻) ions in equilibrium: 2H₂O ⇌ H₃O⁺ + OH⁻, with an ion product constant (K_w) of 1.0 × 10⁻¹⁴ at 25°C, yielding concentrations of approximately 10⁻⁷ mol/L for each ion.[60] [61] This results in a specific conductivity of 0.055 μS/cm for ultrapure water at 25°C, making it a poor conductor compared to solutions with dissolved electrolytes.[60] [62] The presence of impurities, such as dissolved salts, minerals, or acids, introduces additional charge-carrying ions (e.g., Na⁺, Cl⁻), dramatically increasing conductivity; for instance, typical tap water can reach 100–1000 μS/cm depending on ionic content.[63] [61] Electrolysis of water involves the electrolytic decomposition of H₂O into hydrogen and oxygen gases using direct current, following the half-reactions: at the cathode, 2H₂O + 2e⁻ → H₂ + 2OH⁻; at the anode, 2H₂O → O₂ + 4H⁺ + 4e⁻ (or simplified in neutral/alkaline media). The overall reaction is 2H₂O → 2H₂ + O₂, with a theoretical minimum cell potential of 1.23 V derived from the standard Gibbs free energy change (ΔG° = 237.2 kJ/mol at 25°C).[64] [65] In practice, overpotentials at electrodes (typically 0.3–1 V total) and ohmic losses necessitate applied voltages of 1.5–2.0 V or higher, often requiring electrolytes like sulfuric acid or potassium hydroxide to enhance conductivity and reduce resistance.[64] [66] Impurities can catalyze side reactions or degrade electrodes, but controlled electrolysis yields stoichiometric gases (2:1 H₂:O₂ by volume) at efficiencies up to 70–80% in industrial setups.[67]Occurrence in the Universe
Detection and Abundance
Water is detected throughout the universe via spectroscopic techniques that capture its distinct rotational, vibrational, and electronic transitions in emission or absorption. In the radio and millimeter/submillimeter wavelengths, ground-based and space telescopes observe rotational lines of water vapor, including the prominent 22 GHz (1.35 cm) ground-state transition often amplified by maser emission in star-forming regions and outflows. Far-infrared observatories like Herschel have mapped extensive water line forests from ortho- and para-water species, revealing its presence in protostellar envelopes and shocks. Infrared spectroscopy identifies vapor through vibrational bands near 6 μm and ice mantles via broad absorption features at 3 μm and 6 μm on dust grains in molecular clouds. Recent advancements with ALMA enable high-resolution imaging of submillimeter water transitions, such as the 448 GHz line first detected in 2017 toward nearby galaxies, confirming water's role in distant star formation.[68][69] The first detection of the water molecule in the interstellar medium occurred in 1969 through radio observations of absorption against the Sagittarius B2 complex, identifying H₂O via its 22 GHz line. Subsequent surveys have confirmed water in diverse environments, from comets and interstellar objects like 2I/Borisov—where Swift telescope UV observations quantified water production rates—to exoplanet atmospheres via transmission spectroscopy during transits, as with Hubble's identification of water vapor in GJ 9827d in 2024. In extragalactic contexts, ALMA has traced water emission in high-redshift galaxies, such as SPT0311-58 at z=3, billions of light-years away, linking it to molecular gas reservoirs fueling early star formation.[70][71] Regarding abundance, water ranks among the most common molecules after H₂ and CO in the interstellar medium, though its distribution varies sharply by environment. In cold, dense molecular clouds (n_H ≈ 10^4 cm⁻³, T ≈ 10 K), ~90% of water resides as ice on dust grains with abundances relative to total hydrogen of ~10^{-4}, formed via successive hydrogenation of atomic oxygen on grain surfaces. Gaseous water vapor is scarcer in these regions (~10^{-7} relative to H₂) due to freeze-out, but abundances rise to ~10^{-5}–10^{-4} in warmer shocked or irradiated gas, as in outflows or photon-dominated regions. Early cosmic water originated in population III supernovae at z > 20, with simulations showing efficient O + H₂ → OH → H₂O synthesis yielding up to 10^{50} molecules per event. Vast reservoirs exist around quasars, such as the 10^{13} solar masses of water vapor (equivalent to 140 trillion Earth oceans) detected in APM 08279+5255 at z=3.91 via its 658 GHz line. Despite ubiquity, water's fractional abundance on many exoplanets remains low, often <1% in atmospheres, as inferred from JWST and Hubble spectra.[72][73]Forms and Exotic States
Water manifests in numerous phases beyond the familiar liquid, solid (ice Ih), and vapor states, particularly under extreme conditions prevalent in astrophysical environments. Amorphous ice, lacking long-range crystalline order, dominates in the interstellar medium and on cold celestial bodies, formed by vapor deposition at temperatures below 130 K; it exists in low-density (LDA, ~0.94 g/cm³) and high-density (HDA, ~1.17 g/cm³) variants, with recent discoveries of intermediate-density forms resembling liquid water's structure more closely.[74][75] Crystalline polymorphs of ice number over 20, including hexagonal ice Ih stable at ambient pressures, cubic ice Ic in clouds, and high-pressure phases like ice VII (stable above 2.1 GPa) and ice X (above 100 GPa, with symmetric hydrogen bonds). These polymorphs arise from hydrogen bonding arrangements under varying pressure and temperature, influencing planetary interiors and cometary structures.[76] Superionic ice, a hybrid phase where oxygen atoms form a body-centered cubic lattice while hydrogen ions diffuse freely like a liquid, emerges at pressures exceeding 50 GPa and temperatures of 1000–3000 K; experimentally confirmed in diamond anvil cells, it exhibits high electrical conductivity and opacity, potentially constituting a significant fraction of water in Uranus and Neptune's mantles, explaining their anomalous magnetic fields.[77][78][79] Supercritical water, beyond the critical point of 374°C (647 K) and 22 MPa, lacks distinct liquid-vapor boundaries, exhibiting gas-like diffusivity and liquid-like density; this state occurs in deep hydrothermal systems and may influence chemistry in hot, pressurized exoplanetary atmospheres or stellar envelopes.[80][81] Emerging observations include plastic ice VII, a disordered yet rigid phase predicted by simulations and detected experimentally in 2025, blending solid-like mechanics with liquid-like dynamics at gigapascal pressures. Such exotic states underscore water's polymorphism, driven by quantum effects and hydrogen bond networks, with implications for cosmology from Oort cloud ices to giant planet dynamos.[82]Role in Planetary Habitability
Liquid water is considered a prerequisite for habitability on planetary bodies because it enables the chemical reactions necessary for life as known from Earth-based biology, serving as a universal solvent that dissolves a wide range of substances and facilitates metabolic processes in cells. All observed life forms require liquid water to maintain cellular structures, transport ions and molecules, and catalyze enzymatic reactions, with no empirical exceptions documented despite extensive terrestrial and extremophile studies. [83] [84] The stability of liquid water depends on temperature ranges typically between 0°C and 100°C at standard atmospheric pressure, though this can extend under varying pressures, such as in subsurface oceans where hydrostatic pressure prevents freezing. [85] [86] The concept of the habitable zone (HZ) centers on the orbital distance from a star where a planet with an Earth-like atmosphere can sustain surface liquid water, generally spanning from the inner edge where water vaporizes to the outer edge where it freezes, modulated by stellar luminosity and spectral type. For Sun-like stars, this zone extends approximately from 0.95 to 1.67 astronomical units, as calculated from radiative-convective models balancing incoming stellar flux with planetary albedo and greenhouse effects. [85] [87] Factors such as planetary mass, atmospheric composition (e.g., CO₂ or H₂O vapor enhancing greenhouse warming), and internal heat from radiogenic decay or tidal forces can expand habitability beyond the classical HZ, as evidenced by potential subsurface liquid water on moons like Europa, where tidal heating maintains oceans beneath ice shells up to 100 km thick. [88] [86] Water's thermodynamic properties further enhance habitability by buffering environmental fluctuations: its high latent heat of vaporization (2,260 kJ/kg) and specific heat capacity (4.18 J/g·°C) stabilize surface temperatures against diurnal or seasonal variations, while phase transitions drive hydrological cycles that distribute heat and nutrients globally. [89] On water-rich worlds, excessive ocean coverage could limit land-based diversification but still permit habitability if convection and upwelling support nutrient cycling, as modeled for "ocean planets" with depths exceeding 100 km. [90] Empirical searches for exoplanets prioritize HZ candidates with water signatures, such as vapor detected via transmission spectroscopy, underscoring water's role in prioritizing targets for biosignature hunts despite challenges from atmospheric loss or desiccation over billions of years. [91] [92]Hydrology on Earth
Global Distribution
Approximately 1.386 billion cubic kilometers of water exist on Earth, covering about 71 percent of its surface.[4] Of this total volume, saline water constitutes 96.5 percent, primarily in oceans, while freshwater accounts for the remaining 2.5 percent.[4] Oceans dominate the distribution, holding over 97 percent of all water when including minor saline contributions from inland seas and groundwater, with the Pacific Ocean alone comprising roughly half of the oceanic volume at 660 million cubic kilometers.[93] Freshwater is unevenly distributed, with 68.7 percent locked in glaciers and ice caps—predominantly in Antarctica (about 60 percent of global freshwater) and Greenland—rendering much of it inaccessible for immediate human use.[94] Groundwater represents 30.1 percent of freshwater, stored in aquifers beneath continents, though salinity and depth limit usability in many regions.[94] Surface freshwater, including lakes, swamps, and rivers, comprises just 0.3 percent of total freshwater (or 0.009 percent of all water), with Lake Baikal holding the largest single volume at 23,615 cubic kilometers.[94]| Water Type | Percentage of Total Water | Volume (million km³) |
|---|---|---|
| Oceans (saline) | 96.5% | 1,338 |
| Glaciers and ice caps | 1.74% | 24.1 |
| Groundwater (fresh) | 0.76% | 10.5 |
| Surface water (fresh) | 0.013% | 0.18 |
| Atmosphere (vapor) | 0.001% | 0.013 |
| Rivers and biosphere | <0.0001% | Negligible |