Chapter 2

Water

2.1 PERSPECTIVE

As the Earth’s human population continues its exponential increase, the importance of water to the preservation of the standard of living we humans enjoy is becoming of utmost importance. Water is the substance without which we know life, as currently understood, could not exist. The examination of water ranges from the accounting of the vast quantities lying in the oceans and under the surface of the Earth to the minutest details of the structure of water, allowing understanding of its behavior in both natural and contrived systems. As related to environmental process analysis, water is the substance without which there could be no water chemistry. In environmental systems, it is generally water and how water might be affected by a situation or perturbation of a system that drives our desire to understand. Thus, given the importance of water to virtually all that is water chemistry, we will examine important properties of water as related to its structure.

Engineers use many of the physical properties of water in analyses of engineered systems; tables yielding values, correlated with temperature, of density, specific weight, viscosity, surface tension, vapor pressure, and bulk modulus of elasticity are found in most textbooks addressing fluid mechanics. These are mechanical properties but are often important in environmental process analysis. Consideration of the molecular structure and molecular behaviors within liquid water can yield fascinating insights as to why these mechanical properties are as they are. For example, the physical chemists (e.g., Levine, 1988; Williams et al., 1978) tell us that the ordering of the oxygen–hydrogen bonds as water freezes leads to a density of solid water (ice) that is lower than that of liquid water. Consider the alternate existence we would know if the crystallization of water behaved in a manner similar to the crystallization of many other liquids wherein the solid is more dense than the liquid.

The properties of water leading to its rather anomalous behavior relative to other liquids are those that also govern the behavior of water in interactions with solutes—constituents present in and intimately mixed within the water. The term “dissolved” seems to have functional definitions. In the past, we referred to dissolved solids as those not separable from liquid water by a particular glass microfiber filter. In another application, we “filter” sodium and other ions from seawater or brackish water using reverse osmosis. We might use a term like “solvated,” suggesting that the solid somehow has a bond with water in the aqueous solution. It is the particular structure of water that leads to its ability to bond with “solvated” solids. The important properties of water stem from the unique arrangement of electron orbitals around the water molecule. Herein we could launch into a detailed investigation of the quantum chemistry surrounding the water molecule—at which point a typical engineering student’s mind wanders to seemingly more relevant topics. Thus, we will restrict our discussions and associated understandings to the semiquantitative nature.

2.2 IMPORTANT PROPERTIES OF WATER

Based on Pauling’s electronegativity scale (H = 2.2, O = 3.4), we may quite simply understand that hydrogen is quite content to contribute its lone electron to a bond with another atom while oxygen is quite intent upon acquiring two electrons to render its outer electron orbital to be like that of neon, a noble gas. Consequently, each hydrogen atom of a water molecule shares a pair of electrons with the oxygen and two remaining pairs of electrons are largely associated with the oxygen atom. A Lewis dot diagram for water is shown in Figure 2.1. When we consider the three-dimensional nature of the water molecule, the tendency for the electron pairs to orient their molecular orbitals (MOs) as far removed as possible from the other MOs ideally would lead to a tetrahedron as the base shape. Were the structure to be a regular tetrahedron, the H–O–H bond angle would be 109.5 °. Attractions of the shared electron pairs to both the O and an H “thin” the MOs relative to those of the unshared pairs. Then, the unshared electron pairs exert further influence to “push” the MOs of the shared electron pairs closer together. The faces of the tetrahedron are not equilateral triangles. The electrons of the lone pairs exercise greater repulsion on each other, making the lone pair MOs “fatter” than those of the bonded pairs. Further, the lone pair MOs exert greater repulsion on each other than the bonded pair MOs and thus push the bonded pair MOs closer to each other. As a result, the bond angle from the centroid of the hydrogen atom through the centroid of the oxygen atom to the centroid of the other hydrogen atom (H–O–H bond angle) is measured to be 104.5 ° rather than the ideal 109.5 ° (Levine, 1988). In order to visualize the departure from the ideal shape, we set the tetrahedron on the table with the hydrogen atoms and one unshared molecular orbital as the base. A line through the two hydrogen atoms is north–south and the unshared molecular orbital is to the east. The remaining unshared molecular orbital then is at the apex. Then, relative to the apex of a regular tetrahedron, the true apex would be displaced upward and to the west. The north–south line connecting the two hydrogen atoms would be shorter than that of the regular tetrahedron. The west face of the tetrahedron would be an isosceles triangle with a base shorter than the other two sides. The northeast and southeast faces would be isosceles triangles with the side oriented to the east as the longest side. The base would be an isosceles triangle of shape identical to the westward oriented face.

FIGURE 2.1 Lewis “dot” diagram for water.

The electronegativity of the oxygen relative to the hydrogen atoms leads to the well-known polarity of the water molecule. The bonded pair electrons exist in MOs that are associated with both the hydrogen and the oxygen. As a consequence of the greater electronegativity of oxygen, the electrons have a higher probability of residing in a portion of the MO associated with the oxygen atom than with the hydrogen atom. The consequence of this probability is the familiar partial positive (δ+) charges assigned to the hydrogen atoms and partial negative (δ–) charge assigned to the oxygen. The requirement for electroneutrality leads us to conclude that δ– is twice δ+. The positive charge is concentrated at each of the hydrogen atoms and the negative charge is concentrated along the line connecting the centroids of the two nonshared MOs. This concentration of negative charge is responsible for the capability for the bonding of a proton with a water molecule to form the hydronium ion. Were we to allow the centroids of the hydrogen and oxygen atoms to define a plane and to develop a shorthand diagram of the water molecule, we might arrive at something similar to the depiction shown in Figure 2.2.

When we examine this shorthand structure, we may easily understand that hydrogen bonding (interaction between the partial positive of the hydrogen with the partial negative of the oxygen) within liquid water can lead to the formation of a structure within the liquid. Williams et al. (1978) and Stumm and Morgan (1996) refer to “clusters” of structured water molecules within the liquid separated by regions of free, molecular water, shown pictorially in Figure 2.3. Within the clusters, water molecules have a “structure,” with obviously shorter average bond distances than in crystalline ice. At the temperature of its maximum density (3.98 °C) the predominance of these clusters is at maximum. As temperature is raised, the predominance of clusters is decreased until at the boiling point, clustered water is at minimum. As temperature is increased from 3.98 °C, the density of water is decreased as a consequence of the longer hydrogen bonds predominant in the free water. As temperature is reduced below 3.98 °C, the ordering of the hydrogen–oxygen bonds into a structure more like that of crystalline ice renders the solution to be less dense. More detailed discussions of these “clusters” and of their “flickering” nature are presented by Williams et al. (1978) and by various texts addressing water chemistry (e.g., Brezonik and Arnold, 2011; Stumm and Morgan, 1996). The physical chemists have modeled the various properties of water using this structure in combination with the Valence Shell Electron Pair Repulsion (VSEPR) method and attained surprising agreement between model predictions and experimental observations (Levine, 1988). We will leave such endeavors to the physical and quantum chemists. Herein, we are much more interested in understanding the manifestations of these subatomic properties on the interactions of water molecules with solutes residing within the liquid water.

FIGURE 2.2 Shorthand structure for the water molecule.

FIGURE 2.3 (a) Hydrogen-bonded open tetrahedral structure of ice. (b) Frank–Wen flickering cluster model of liquid water. Reproduced from Stumm and Morgan (1996) with permission from John Wiley & Sons.

Of particular interest are the interactions between water and charged entities—ions—within an aqueous solution. The partial negative of the oxygen tends to orient with the positive...