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The self-ionization of water (also autoionization of water, and autodissociation of water, or simply dissociation of water) is an ionization reaction in pure water or in an aqueous solution, in which a water molecule, H2O, deprotonates (loses the nucleus of one of its hydrogen atoms) to become a hydroxide ion, OH−. The hydrogen nucleus, H+, immediately protonates another water molecule to form a hydronium cation, H3O+. It is an example of autoprotolysis, and exemplifies the amphoteric nature of water.
History and notation
The self-ionization of water was first proposed in 1884 by Svante Arrhenius as part of the theory of ionic dissociation which he proposed to explain the conductivity of electrolytes including water. Arrhenius wrote the self-ionization as . At that time, nothing was yet known of atomic structure or subatomic particles, so he had no reason to consider the formation of an ion from a hydrogen atom on electrolysis as any less likely than, say, the formation of a ion from a sodium atom.
In 1923 Johannes Nicolaus Brønsted and Martin Lowry proposed that the self-ionization of water actually involves two water molecules: . By this time the electron and the nucleus had been discovered and Rutherford had shown that a nucleus is very much smaller than an atom. This would include a bare ion which would correspond to a proton with zero electrons. Brønsted and Lowry proposed that this ion does not exist free in solution, but always attaches itself to a water (or other solvent) molecule to form the hydronium ion (or other protonated solvent).
Later spectroscopic evidence has shown that many protons are actually hydrated by more than one water molecule. The most descriptive notation for the hydrated ion is , where aq (for aqueous) indicates an indefinite or variable number of water molecules. However the notations and are still also used extensively because of their historical importance. This article mostly represents the hydrated proton as , corresponding to hydration by a single water molecule.
Equilibrium constant
Chemically pure water has an electrical conductivity of 0.055 μS/cm. According to the theories of Svante Arrhenius, this must be due to the presence of ions. The ions are produced by the water self-ionization reaction, which applies to pure water and any aqueous solution:
- H2O + H2O ⇌ H3O+ + OH−
Expressed with chemical activities a, instead of concentrations, the thermodynamic equilibrium constant for the water ionization reaction is:
which is numerically equal to the more traditional thermodynamic equilibrium constant written as:
under the assumption that the sum of the chemical potentials of H+ and H3O+ is formally equal to twice the chemical potential of H2O at the same temperature and pressure.[1]
Because most acid–base solutions are typically very dilute, the activity of water is generally approximated as being equal to unity, which allows the ionic product of water to be expressed as:[2]
In dilute aqueous solutions, the activities of solutes (dissolved species such as ions) are approximately equal to their concentrations. Thus, the ionization constant, dissociation constant, self-ionization constant, water ion-product constant or ionic product of water, symbolized by Kw, may be given by:
where is the molarity (molar concentration)[3] of hydrogen cation or hydronium ion, and is the concentration of hydroxide ion. When the equilibrium constant is written as a product of concentrations (as opposed to activities) it is necessary to make corrections to the value of depending on ionic strength and other factors (see below).[4]
At 24.87 °C and zero ionic strength, Kw is equal to 1.0×10−14. Note that as with all equilibrium constants, the result is dimensionless because the concentration is in fact a concentration relative to the standard state, which for H+ and OH− are both defined to be 1 molal (= 1 mol/kg) when molality is used or 1 molar (= 1 mol/L) when molar concentration is used. For many practical purposes, the molality (mol solute/kg water) and molar (mol solute/L solution) concentrations can be considered as nearly equal at ambient temperature and pressure if the solution density remains close to one (i.e., sufficiently diluted solutions and negligible effect of temperature changes). The main advantage of the molal concentration unit (mol/kg water) is to result in stable and robust concentration values which are independent of the solution density and volume changes (density depending on the water salinity (ionic strength), temperature and pressure); therefore, molality is the preferred unit used in thermodynamic calculations or in precise or less-usual conditions, e.g., for seawater with a density significantly different from that of pure water,[3] or at elevated temperatures, like those prevailing in thermal power plants.
We can also define pKw −log10 Kw (which is approximately 14 at 25 °C). This is analogous to the notations pH and pKa for an acid dissociation constant, where the symbol p denotes a cologarithm. The logarithmic form of the equilibrium constant equation is pKw = pH + pOH.
Dependence on temperature, pressure and ionic strength
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The dependence of the water ionization on temperature and pressure has been investigated thoroughly.[5] The value of pKw decreases as temperature increases from the melting point of ice to a minimum at c. 250 °C, after which it increases up to the critical point of water c. 374 °C. It decreases with increasing pressure.
| Temperature | Pressure[7] | pKw |
|---|---|---|
| 0 °C | 0.10 MPa | 14.95 |
| 25 °C | 0.10 MPa | 13.99 |
| 50 °C | 0.10 MPa | 13.26 |
| 75 °C | 0.10 MPa | 12.70 |
| 100 °C | 0.10 MPa | 12.25 |
| 150 °C | 0.47 MPa | 11.64 |
| 200 °C | 1.5 MPa | 11.31 |
| 250 °C | 4.0 MPa | 11.20 |
| 300 °C | 8.7 MPa | 11.34 |
| 350 °C | 17 MPa | 11.92 |
With electrolyte solutions, the value of pKw is dependent on ionic strength of the electrolyte. Values for sodium chloride are typical for a 1:1 electrolyte. With 1:2 electrolytes, MX2, pKw decreases with increasing ionic strength.[8]
The value of Kw is usually of interest in the liquid phase. Example values for superheated steam (gas) and supercritical water fluid are given in the table.
Comparison of pKw values for liquid water, superheated steam, and supercritical water.[1] Temp.Pressure350 °C 400 °C 450 °C 500 °C 600 °C 800 °C 0.1 MPa 47.961b 47.873b 47.638b 46.384b 40.785b 17 MPa 11.920 (liquid)a 25 MPa 11.551 (liquid)c 16.566 18.135 18.758 19.425 20.113 100 MPa 10.600 (liquid)c 10.744 11.005 11.381 12.296 13.544 1000 MPa 8.311 (liquid)c 8.178 Zdroj:https://en.wikipedia.org?pojem=Self-ionization_of_water
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