15.3.1 Ice
Solidification. When pure water is cooled to below 0 8 C and heat is
continuously removed, all of the water will crystallize (assuming for the
moment that ice nucleation occurs readily). This is a case ofcrystallization
from the melt, often called solidification. Crystallization is a cooperative
transition and it completely occurs at the equilibrium temperature, i.e., 0 8 C.
That means that we cannot speak of a supersaturation. Proceeding from Eq.
(14.3), we derive the driving force for crystallization, i.e., the difference in
chemical potential between ice and water, as
Dm¼DL!SG¼DL!SHð 1 T
TeqÞ& 22 ðTeqTÞð 15 : 7 Þwhere the part after the&sign is valid for ice formation, if expressed in SI
units. It is assumed thatDHandDSare independent of temperature, which
is not precisely true. See Appendix I for further data.
In all foods, the water contains solutes, but the ice formed will have
the same crystal structure as that formed from pure water.
Crystal Structure. Water can solidify in eight different stable
modifications or enantiotropic polymorphs. Seven of these form at very high
pressures—generally above 200 MPa—and need not be considered here. The
ice formed under common conditions is called ice I; the crystal system is
hexagonal. The unit cell contains four water molecules; these are for the
most part kept together by hydrogen bonds, leading to a fairly open
structure, as is reflected in the low density. The other crystal modifications
have far higher densities.
Hydrogen bonds are fairly strong; according to Table 3.1 they are of
the order of 10RTor 22 kJ per mole. In ice, every molecule is H-bonded to
four neighbors, in water to about three; this would mean a decrease of 0.5
bond per molecule, corresponding to about 11 kJ?mol^1. This is of the same
order as the enthalpy of fusion: 6 kJ?mol^1. The latter value is high, and
also the melting temperature of ice is high, for such a small molecule. The
relatively open structure (low density) of ice is also due to hydrogen
bonding. The energy of these bonds is strongly dependent on bond angle,
and to minimize the free energy the molecules in ice are so oriented that the
bonds (O—HO) are almost linear.
‘‘Anomalies.’’ Figure 15.13 gives some properties of water at low
temperature. It is seen that water has its maximum density at 4 8 C, and that
also below 0 8 C density keeps decreasing with decreasing temperature. The