Hydrogen Bonding in Minerals

Hydrogen bonding is a weak electrostatic attraction found in many minerals, arising from the unequal distribution of electrical charge within polar molecules or crystal layers. ^[1] It is not a primary valence bond-it does not involve the sharing or transfer of electrons in the way ionic, covalent, or metallic bonding does. Instead, hydrogen bonds act between molecules or between the layers of a crystal, linking structural units that are already internally stable through stronger bonds. Because they are weak, hydrogen bonds have a predictable effect on mineral properties: minerals held together partly or entirely by hydrogen bonding tend to be soft, easily cleaved, and structurally responsive to changes in temperature.

The Polarity Mechanism

The origin of hydrogen bonding lies in the electrical polarity created when hydrogen forms a bond with a strongly electronegative element such as oxygen. In a water molecule, each of the two hydrogen atoms shares its single electron with one of the oxygen’s 2p orbitals, forming a dominantly covalent bond. ^[1] Because the electronegativity difference between oxygen and hydrogen is 1.4, oxygen has a much stronger claim on the shared electrons than hydrogen does. ^[1] Electron density shifts toward the oxygen, leaving the hydrogen nuclei exposed as regions of net positive charge, while the far side of the oxygen acquires a net negative charge. This produces an electrical polarity to the water molecule (Figure 3.13b)-positive near the two H nuclei and negative at two nodes on the opposite side of the O. ^[1] The four charge concentrations-two positive at the hydrogen positions, two negative on the oxygen side-point toward the corners of a tetrahedron, giving the molecule a well-defined three-dimensional polarity. ^[1] This tetrahedral arrangement of charge is what controls the geometry of hydrogen-bonded structures, because each positive charge site on one molecule is attracted to a negative charge site on a neighboring molecule.

Ice: A Mineral Built from Hydrogen Bonds

Ice is the most familiar mineral in which hydrogen bonding is prominent. It is composed of molecules of H₂O bonded together to form snow, hail, glaciers, and ice. ^[1] Below 0°C, the thermal energy of the molecules drops low enough that the weak electrostatic attraction between the positive hydrogen end of one molecule and the negative oxygen end of a neighbor can hold the molecules in fixed positions. ^[1] The result is a hexagonal framework in which every water molecule bonds to exactly four neighbors through hydrogen bonds-matching the four corners of the tetrahedral charge distribution described above. ^[1] This hexagonal framework is directly visible at the macroscopic scale in the six-fold symmetry of snowflakes: the molecular geometry of hydrogen bonding propagates outward through the crystal to produce a six-sided crystal habit.

Serpentine: Hydrogen Bonds Between Layers

In sheet silicates, hydrogen bonding links adjacent structural layers together rather than holding molecules in a framework. Serpentine is the most important example. Its structure consists of stacked layers, each built from three planes of anions arranged in order from bottom to top: pure oxygen (O²⁻), a mixed plane of oxygen and hydroxyl (OH⁻), and pure hydroxyl at the top. ^[1] Within each layer, Mg²⁺ and Si⁴⁺ cations fill the interstices between the anion planes, forming bonds that satisfy all valence requirements internally. ^[1] Bonding between layers is then accomplished by the attraction between the positively charged hydrogen cations of the OH⁻ groups on the top surface of one layer and the negatively charged O²⁻ anions on the bottom of the layer above it. ^[1] Because these interlayer hydrogen bonds are weaker than the ionic and covalent bonds within each layer, serpentine cleaves very readily parallel to the layers-it takes much less energy to pry apart the layers by breaking hydrogen bonds than it would to cut through the layer itself.

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