Chemistry Of Water

Water molecules (H2O) are symmetric (point group C2ν) with two mirror planes of symmetry and a 2-fold rotation axis.
The hydrogen atoms may possess parallel (ortho-H2O) or antiparallel (para-H2O) nuclear spin.

The equilibrium ratio is all para at zero Kelvin shifting to 3:1 ortho: para at less cold temperatures (>50 K); the equilibrium taking months to establish itself in ice and nearly an hour in ambient water.

This means that liquid H2O effectively consists of a mixture of non-identical molecules.

Many materials preferentially adsorb para-H2O due to its non-rotation ground state.

The apparent difference in energy between the two states is a significant 1-2 kJ mol-1, far greater than expected from spin-spin interactions (< μJ mol-1).

Early 5-point molecular models, with explicit negative charge where the lone pairs are purported to be, faired poorly in describing hydrogen bonding, but a recent TIP5P model shows some promise.

Although there is no apparent consensus of opinion, such descriptions of substantial sp3-hybridized lone pairs in the isolated water molecule should perhaps be avoided, as an sp2-hybridized structure (plus a Pz orbital) is indicated. This rationalizes the formation of (almost planar) trigonal hydrogen bonding that can be found around some restricted sites in the hydration of proteins and where the numbers of hydrogen bond donors and acceptors are unequal.

The electron density distribution for water is shown above right with some higher density contours around the oxygen atom omitted for clarity. The polarizability of the molecule is centered around the O-atom (1.4146 Å3) with only small polarizabilities centered on the H-atoms (0.0836 Å3). For an isolated H216O, H217O or H218O molecule, the calculated O-H length is 0.957854 Å and the H-O-H angle is 104.500° (D216O, 0.957835 Å, 104.490°).

The charge distribution depends significantly on the atomic geometry and the method for its calculation but is likely to be about -0.7e on the O-atom (with the equal but opposite positive charge equally divided between the H-atoms) for the isolated molecule. The experimental values for gaseous water molecules are O-H length 0.95718 Å, H-O-H angle 104.474°.

These values are not maintained in liquid water, where ab initio (O-H length 0.991 Å, H-O-H angle 105.5° and neutron diffraction studies (O-H length 0.970 Å, H-O-H angle 106° suggest slightly greater values, which are caused by the hydrogen bonding weakening the covalent bonding.

These bond lengths and angles are likely to change, due to polarization shifts, in different hydrogen-bonded environments and when the water molecules are bound to solutes and ions.

Commonly used molecular models utilize O-H lengths of between 0.957 Å and 1.00 Å and H-O-H angles of 104.52° to 109.5°.

The electronic structure has been proposed as 1sO2.00 2sO1.82 2pxO1.50 2pzO1.12 2pyO2.00 1sH10.78 1sH20.78, however, it now appears that the 2s orbital may be effectively unhybridized with the bond angle expanded from the (then) expected angle of 90° due to the steric and ionic repulsion between the partially-positively charged hydrogen atoms (as proposed by Pauling over 50 years ago.

The molecular orbitals of water, (1a1)2(2a1)2(1b2)2(3a1)2(1b1)2, are shown on another page (24 KB).

Shown opposite is the electrostatic potential associated with the water structure. Although the lone pairs of electrons do not appear to give distinct directed electron density in isolated.

The mean van der Waals diameter of water has been reported as identical to that of isoelectronic neon (2.82 Å) .

Molecular model values and intermediate peak radial distribution data indicate however that it is somewhat greater (~3.2Å). The molecule is clearly not spherical, however, with about a ±5% variation in van der Waals diameter dependent on the axis chosen; approximately tetrahedrally placed slight indentations being apparent opposite the (putative) electron pairs.

Much effort has been expended on the structure of small isolated water clusters. The most energetically favorable water dimer is shown below with a section through the electron density distribution (high densities around the oxygen atoms have been omitted for clarity).

This shows the tetrahedrally of the bonding in spite of the lack of clearly seen lone pair electrons; although a small amount of distortion along the hydrogen bond can be seen. This tetrahedrally is primarily caused by electrostatic effects (i.e. repulsion between the positively charged non-bonded hydrogen atoms) rather than the presence of tetrahedrally placed lone pair electrons.

The hydrogen-bonded proton has reduced electron density relative to the other protons.

Note that, even at temperatures as low as a few kelvins, there are considerable oscillations (< ps) in the hydrogen bond length and angles. The molecular orbitals of the water dimer are shown on another page (50 KB)R = 2.976 (+0.000, -0.030) Å, α = 6 ± 20°, β = 57 ± 10°; α is the donor angle and β is the acceptor angle.

The dimer (with slightly different geometry) dipole moment is 2.6 D.

Although β is close to as expected if the lone pair electrons were tetrahedrally placed (= 109.47°/2), the energy minimum (~21 kJ mol-1) is broad and extends towards β = 0°.

Simplified models for the water molecule have been developed to agree with particular physical properties (e.g. agreement with the critical parameters) but they are not robust and resultant data are often very sensitive to the precise model parameters [206].

Models are still being developed and are generally more complex than earlier but they still generally have poor predictive value outside the conditions and physical parameters for which they were developed.

Although not often perceived as such, water is a very reactive molecule available at a high concentration. This reactivity, however, is greatly moderated at ambient temperatures due to the extensive hydrogen bonding.

Water molecules each possess a strongly nucleophilic oxygen atom that enables many of life’s reactions, as well as ionizing to produce reactive hydrogen and hydroxide ions.

Reduction of the hydrogen bonding at high temperatures, or due to electromagnetic fields, results in greater reactivity of the water molecules.

Water’s composition (two parts hydrogen to one part oxygen) was discovered by the London scientist Henry Cavendish (1731-1810) in about 1781. He reported his findings in terms of phlogiston (later the gas he made was proven to be hydrogen) and dephlogisticated air (later this was proven to be oxygen). Cavendish died (1810) in his Laboratory just 30 minutes walk from the present site of London South Bank University.

It has recently been suggested that H1.5O may better reflect the formula at very small (attosecond) timescales when some of the H-atoms appear invisible to neutron and electron interaction.

These results have been questioned and are now thought erroneous, but such suggestions do add support to the view that observations concerning the structure of water should be tempered by the timescale used.

The tetrahedral angle is 180-cos-1(1/3)°; 109.47122° = 109° 28′ 16.39″. Tetrahedral (q, the orientational order parameter) may be defined as, where φjk is the angle formed by lines drawn between the oxygen atoms of the four nearest and hydrogen-bonded water molecules.

It equals unity for perfectly tetrahedral bonding (where cos(φjk) = -1/3) and averages zero (±0.5 SD) for random arrangements, with a minimum value of -3.

Due to the deuterium’s nuclear spin of 1 (cf. 1/2 for H’s spin), the lowest energy form of D2O is ortho. D2O converts to a 2:1 ortho: para ratio at higher temperatures.

HDO, having non-equivalent hydrogen atoms, does not possess an ortho/para distinction. T2O behaves similarly to H2O as tritium also possesses a nuclear spin of 1/2.

The charge on the hydrogen atoms across the periodic table is shown opposite. The hydrogen atom charges are blue and the charges on the other atoms are indicated red.

The actual values depend on the vibrational state of the molecule with even values of 180° being attainable during high order bend vibrations (v2 >= 7, λ < 900 nm) for the H-O-H angle.

Vibrations are asymmetric around the mean positions. In the ground state, the bond angle (104.5°) is much closer to the tetrahedral angle than that of the other Group VI hydrides, H2S (92.1°), H2Se (91°), or H2Te (89°).

The H-O-H angle in ice Ih is reported as 106.6°±1.5°, whereas recent modeling gives values of 108.4°±0.2° for ice Ih and 106.3°±4.9° for water.

The polarity of water
Water has a simple molecular structure, It is composed of one oxygen atom and two hydrogen atoms.

Each hydrogen atom is covalently bonded to oxygen via a shared pair of electrons, Oxygen also has two unshared pairs of electrons.

Thus there are 4 pairs of electrons surrounding the oxygen atom, two pairs involved in covalent bonds with hydrogen, and two unshared pairs on the opposite side of the oxygen atom., Oxygen is an “electronegative” or electron “loving” atom compared with hydrogen.

Water is a “polar” molecule, meaning that there is an uneven distribution of electron density.

Water has a partial negative charge ( ) near the oxygen atom due to the unshared pairs of electrons, and partial positive charges ( ) near the hydrogen atoms.
An electrostatic attraction between the partial positive charge near the hydrogen atoms and the partial negative charge near the oxygen results in the formation of a hydrogen bond as shown in the illustration.

And where the numbers of hydrogen bond donors and acceptors are unequal.

The ability of ions and other molecules to dissolve in water is due to polarity. For example, in the illustration below sodium chloride is shown in its crystalline form and dissolved in water.

 

Many other unique properties of water are due to the hydrogen bonds.

For example, ice floats because hydrogen bonds hold water molecules further apart in a solid than in a liquid, where there is one less hydrogen bond per molecule.

The unique physical properties, including a high heat of vaporization, strong surface tension, high specific heat, and nearly universal solvent properties of water are also due to hydrogen bonding.

The hydrophobic effect or the exclusion of compounds containing carbon and hydrogen (non-polar compounds) is another unique property of water caused by hydrogen bonds.

The hydrophobic effect is particularly important in the formation of cell membranes.

The best description is to say that water “squeezes” non-polar molecules together.

Acids and Bases, Ionization of Water