What is self-ionization of water


Ionization reaction in pure water or in an aqueous solution

The Self-ionization of water (Likewise Autoionization of water, and Autodissociation of water) is an ionization reaction in pure water or in an aqueous solution in which a water molecule, H.2O, deprotonates (loses the nucleus of one of its hydrogen atoms) and becomes a hydroxide ion, OH– –. The hydrogen nucleus H.+immediately protonates another water molecule to form hydronium, H.3Ö+. It is an example of autoprotolysis and illustrates the amphoteric nature of water.

Equilibrium constant [edit]

Animation of the self-ionization of water

Chemically pure water has an electrical conductivity of 0.055 μS / cm. According to Svante Arrhenius' theories, this must be due to the presence of ions. The ions are generated by the self-ionization reaction of water, which applies to pure water and every aqueous solution:

H.2O + H.2O ⇌ H.3Ö+ + OH– –

In terms of chemical activities, instead of concentrations, the thermodynamic equilibrium constant for the water ionization reaction is:

This is numerically equal to the more traditional thermodynamic equilibrium constant, written as:

assuming that the sum of the chemical potentials of H.+ and H.3Ö+ is formally equal to twice the chemical potential of H.2O at the same temperature and pressure.[1]

Since most acid-base solutions are typically very dilute, the activity of water is generally approximated as equal to one, whereby the ionic product of water can be expressed as:[2]

In dilute aqueous solutions, the activities of the dissolved particles are approximately equal to their concentrations. And so it happened that the Ionization constant, Dissociation constant, Self-ionization constant, Water ion product constant or ionic product Water, symbolized by K.w, can be given by:

where [H3O+] is the molarity (≈ molar concentration)[3] of hydrogen or hydronium ions and [OH] is the concentration of the hydroxide ion. If the equilibrium constant is written as the product of concentrations (as opposed to activities), corrections must be made to the value of

depending on the ionic strength and other factors (see below).[4]

At 25 ° C and an ionic strength of zero K.w corresponds to 1.0 × 10−14. Note that, as with all equilibrium constants, the result is dimensionless because the concentration is actually a concentration relative to the standard state that is required for H.+ and oh– – are both defined as 1 mole (or mole). For most practical purposes, the molar and molar concentrations are the same near ambient temperature and pressure. The molar concentration scale leads to concentration values ​​that take into account changes in density with changes in temperature or pressure; Hence, it is the scale that is used in precise or non-ambient applications, e.g. B. for sea water.[3] or at elevated temperatures such as in thermal power plants.

We can also define pK.w

−log10K.w (that's about 14 at 25 ° C). This is analogous to the terms pH and pK.a for an acid dissociation constant, where the symbol p denotes a logarithm. The logarithmic form of the equilibrium constant equation is pK.w = pH + pOH.

Dependence on temperature, pressure and ionic strength [edit]

Temperature dependence of the water ionization constant at 25 MPa
Pressure dependence of the water ionization constant at 25 ° C.
Variation of pK.w with ionic strength of NaCl solutions at 25 ° C.

The dependence of water ionization on temperature and pressure was examined in detail.[5] The value of pK.w decreases with increasing temperature from the melting point of the ice to a minimum at c. 250 ° C, then it rises to the critical point of the water c. 374 ° C. It decreases with increasing pressure.

temperature print[7]pK.w
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

In the case of electrolyte solutions, the value of pK.w depends on the ionic strength of the electrolyte. Sodium chloride values ​​are typical for a 1: 1 electrolyte. With 1: 2 electrolytes, MX2, pK.w decreases with increasing ionic strength.[8]

The value of K.w is usually of interest in the liquid phase. Example values ​​for superheated steam (gas) and supercritical water fluid are given in the table.



350 ° C.400 ° C.450 ° C.500 ° C.600 ° C.800 ° C.
0.1 MPa 47.961b47.873b47.638b46,384b40,785b
17 MPa 11,920 (liquid)a
25 MPa 11,551 (liquid)c16.56618.13518.75819.42520.113
100 MPa 10,600 (liquid)c10.74411.00511.38112.29613.544
1000 MPa 8,311 (liquid)c8.1788.0848.0197,9527,957
Notes on the table. The values ​​apply to supercritical fluids with the exception of the specified: a at saturation pressure corresponding to 350 ° C. b superheated steam. ccompressed or supercooled liquid.

Isotope effects [edit]

Heavy water, D.2O, ionizes less by itself than normal water, H.2Ö;

D.2O + D.2O ⇌ D.3Ö+ + OD– –

This is due to the equilibrium isotope effect, a quantum mechanical effect that is due to the fact that oxygen forms a somewhat stronger bond with deuterium, since the larger deuterium mass leads to a lower zero-point energy.

Expressed with activities aInstead of concentrations, the thermodynamic equilibrium constant for the heavy water ionization reaction is:

Assuming the activity of the D.2O 1 and assume that the D.3Ö+ and OD– – are closely approximated by their concentrations

The following table compares the values ​​of pK.w for H2O and D.2Ö.[9]

T / ° C. 102025304050
H.2Ö 14.53514.16713.99713.83013.53513.262
D.2Ö 15.43915.04914.86914.69914.38514.103

Ionization equilibria in water-heavy water mixtures [edit]

Several species are involved in equilibria between water and heavy water: H.2O, HDO, D.2OH3Ö+, D.3Ö+, H.2DO+, HD2Ö+, HO– –, DO– –.

Mechanism [edit]

The reaction rate for the ionization reaction

2 H.2O → H.3Ö+ + OH– –

depends on the activation energy, ΔE.. According to the Boltzmann distribution, the proportion of water molecules that have sufficient energy due to the thermal population is given by

Where k is the Boltzmann constant. Thus, a certain dissociation can occur, since sufficient thermal energy is available. The following sequence of events has been proposed based on variations in the electric field in liquid water.[10] Occasionally random fluctuations in molecular movements (approximately every 10 hours per water molecule[11]) create an electric field strong enough to break an oxygen-hydrogen bond, resulting in a hydroxide (OH)– –) and hydronium ion (H.3Ö+); The hydrogen nucleus of the hydronium ion migrates along water molecules by the Grotthuss mechanism, and a change in the hydrogen bonding network in the solvent isolates the two ions, which are stabilized by solvation. However, within 1 picosecond, a second reorganization of the hydrogen bonding network enables rapid proton transfer along the electrical potential difference and subsequent recombination of the ions. This timescale corresponds to the time it takes hydrogen bonds to reorient in water.[12][13][14]

The inverse recombination reaction

H.3Ö+ + OH– – → 2 H.2Ö

is one of the fastest known chemical reactions with a reaction rate constant of 1.3 × 1011 M.−1 s−1 at room temperature. Such a rapid rate is characteristic of a diffusion-controlled reaction, in which the rate is limited by the rate of molecular diffusion.[15]

Relationship to the neutral point of the water [edit]

Water molecules dissociate in equal amounts from H.3Ö+ and oh– –, so their concentrations are equal to 1.00 × 10−7 mol dm−3 at 25 ° C. A solution in which the H.3Ö+ and oh– – Concentrations that are equal to each other are considered to be a neutral Solution. In general, the pH of the neutral point is numerically equal to 1 / .2pK.w.

Pure water is neutral, but most water samples contain impurities. If an impurity is an acid or base, it affects the concentrations of hydronium ions and hydroxide ions. Water samples exposed to air absorb some carbon dioxide to form carbonic acid (H.2CO3) and the concentration of H.3Ö+ will increase due to the reaction H.2CO3 + H.2O = HCO3– – + H.3Ö+. The concentration of OH– – will decrease so that the product [H3O+][OH] remains constant for a fixed temperature and pressure. This means that these water samples are slightly acidic. If a pH of exactly 7.0 is required, this must be maintained with a suitable buffer solution.

See also [edit]

References [edit]

  1. ^ ab“Release on the ionization constant of H.2Ö ”(PDF). Lucerne: The International Association for the Properties of Water and Steam. August 2007.
  2. ^IUPAC, Compendium of Chemical Terminology, 2nd edition (the “Goldbuch”) (1997). Version corrected online: (2006–) “Autoprotolysis constant”. doi: 10.1351 / goldbook.A00532
  3. ^ abStumm, Werner; Morgan, James (1996). Aquatic chemistry. Chemical equilibria and rates in natural waters (3rd ed.). John Wiley & Sons, Inc. ISBN.
  4. ^Harned, HS; Owen, BB (1958). The physical chemistry of electrolytic solutions (3rd ed.). New York: Reinhold. pp. 635.
  5. ^International Association for the Properties of Water and Steam (IAPWS)
  6. ^Bandura, Andrei V .; Lvov, Serguei N. (2006). “The ionization constant of water over wide ranges of temperature and density” (PDF). Journal of Physical and Chemical Reference Data. 35 (1): 15-30. Bibcode: 2006JPCRD..35… 15B. doi: 10.1063 / 1.1928231.
  7. ^0.1 MPa for T.
  8. ^Harned, HS; Owen, BB (1958). The physical chemistry of electrolytic solutions (3rd ed.). New York: Reinhold. pp. 634-649, 752-754.
  9. ^Lide, DR, ed. (1990). CRC Handbook of Chemistry and Physics (70th Edition). Boca Raton (FL): CRC press.
  10. ^Geissler, PL; Dellago, C .; Chandler, D .; Hutter, J .; Parrinello, M. (2001). “Autoionization in Liquid Water”. science. 291 (5511): 2121-2124. Bibcode: 2001Sci… 291.2121G. CiteSeerX doi: 10.1126 / science.1056991. PMID 11251111.
  11. ^Eigen, M .; De Maeyer, L. (1955). “Investigations on the kinetics of neutralization I”. Z. Elektrochem. 59: 986.
  12. ^Stillinger, FH (1975). Theory and molecular models for water. Adv. Chem. Phys. Advances in Chemical Physics. 31. Pp. 1-101. doi: 10.1002 / 9780470143834.ch1. ISBN.
  13. ^Rapaport, DC (1983). “Hydrogen Bridges in Water”. Mol. Phys.50 (5): 1151-1162. Bibcode: 1983MolPh..50.1151R. doi: 10.1080 / 00268978300102931.
  14. ^Chen, S.-H .; Teixeira, J. (1986). Structure and dynamics of low temperature water according to scattering techniques. Adv. Chem. Phys. Advances in Chemical Physics. 64. Pp. 1-45. doi: 10.1002 / 9780470142882.ch1. ISBN.
  15. ^Tinoco, I .; Sauer, K .; Wang, JC (1995). Physical Chemistry: Fundamentals and Applications in Life Sciences (3rd ed.). Prentice Hall. p. 386.

External links [edit]