Food Biochemistry and Food Processing (2 edition)

BLBS102-c05 BLBS102-Simpson March 21, 2012 12:2 Trim: 276mm X 219mm Printer Name: Yet to Come

5 Water Chemistry and Biochemistry 89

Table 5.1.Properties of Liquid Water at 298 K

Heat of formation
Hf 285.89 kJ mol−^1

Density at 3.98◦C 1.000 g cm^3

Density at 25◦C 0.9970480 g cm^3

Heat capacity 4.17856 J g−^1 K−^1


Hvaporization 55.71 kJ mol−^1

Dielectric constant 80

Dipole moment 6.24× 10 −^30 Cm

Viscosity 0.8949 mPa s

Velocity of sound 1496.3 m s

Volumetric thermal expansion

coefficient

0.0035 cm^3 g−^1 K−^1

mp 195 K, bp 240 K) and HF (molar mass 20, mp 190 K,

bp 293 K). If we compare the hydrogen compounds of elements

from the same group (O, S, Se, and Te), the normal bp of H 2 O

(373 K) is by far the highest among H 2 S, H 2 Se, and H 2 Te.

Much energy (21 kJ mol−^1 ) is required to break the hydrogen

bonds. The strong hydrogen bonds among water molecules in

condensed phase result in anomalous properties, including the

high enthalpies (energies) of fusion, sublimation, and evapora-

tion given in Table 5.1. Internal energies and entropies are also

high.

Densities of water and ice are also anomalous. Ice at 273 K

is 9% less dense than water, but solids of most substances are

denser than their liquids. Thus, ice floats on water, extending 9%

of its volume above water. Water is the densest at 277 K (4◦C).

Being less dense at the freezing point, still water freezes from the

top down, leaving a livable environment for aquatic organisms.

The hydrogen bonding and polarity also lead to aberrant high

surface tension, dielectric constant, and viscosity.

We illustrate the phase transitions between ice, liquid (wa-

ter), and vapor in a phase diagram, which actually shows the

equilibria among the common phases. Experiments under high

pressure observed at least 13 different ices, a few types of amor-

phous solid water, and even the suggestion of two forms of liq-

uid water (Klug 2002, Petrenko and Whitworth 1999). If these

phases were included, the phase diagram for water would be

very complicated.

Solid H 2 O

At 273.16 K, ice, liquid H 2 O, and H 2 O vapor at 611.15 Pa coex-

ist and are at equilibrium; the temperature and pressure define

thetriple-pointof water. At the normal pressure of 101.3 kPa

(1 atm), ice melts at 273.15 K. The temperature for the equilib-

rium water vapor pressure of 101.3 kPa is the bp, 373.15 K.

Under ambient pressure, ice often does not begin to form until

it is colder than 273.15 K, and this is known assupercooling,

especially for ultrapure water. The degree of supercooling de-

pends on volume, purity, disturbances, the presence of dust, the

smoothness of the container surface, and similar factors. Crys-

tallization starts bynucleation, that is, formation of ice-structure

clusters sufficiently large that they begin to grow and become

crystals. Once ice begins to form, the temperature will return to

the freezing point. At 234 K (− 39 ◦C), tiny drops of ultrapure

water would suddenly freeze, and this is known ashomogeneous

nucleation(Franks et al. 1987). Dust particles and roughness of

the surface promote nucleation and help reduce supercooling for

ice and frost formation.

At ambient conditions, hexagonal ice (Ih) is formed.

Snowflakes exhibit the hexagonal symmetry. Their crystal struc-

ture is well known (Kamb 1972). Every oxygen atom has four hy-

drogen bonds around it, two formed by donating its two H atoms,

and two by accepting the H atoms of neighboring molecules. The

hydrogen bonds connecting O atoms are shown in Figure 5.5.

In normal ice, Ih, the positions of the H atoms are random or

disordered. The hydrogen bonds O H O may be slightly bent,

leaving the H O H angle closer to 105◦than to 109.5◦,the

ideal angle for a perfect tetrahedral arrangement. Bending the

hydrogen bond requires less energy than opening the H O H

angle. Bending of the hydrogen bond and the exchange of H

atoms among molecules, forming H 3 O+and OH−in the solids,

Figure 5.5.The crystal structures of ice hexagonalice (Ih) and cubic ice (Ic). Oxygen atoms are placed in two rings in each to point out their

subtle difference. Each line represents a hydrogen bond O H O, and the H atoms are randomly distributed such that on average, every

oxygen atom has two O H bonds of 100 pm. The O H O distance is 275 pm. The idealized tetrahedral bond angles around oxygen are

1095 ◦.