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1 Chpt. 3: Properties of Water and Seawater James W. Murray (9/30

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1 Chpt. 3: Properties of Water and Seawater James W. Murray (9/30 Chpt. 3: Properties of Water and Seawater James W. Murray (9/30/04) Univ. Washington I. The Nature of pure water Seawater is composed mostly of water (H2O). In fact it is about 96.5 wt % water. Sediments are also mostly water. Most fine grained surface sediments have a porosity ( Φ = volume of pores to volume of solids) of greater than 90%. Almost every process we discuss will occur and be affected by water. Thus, water has been called the universal solvent. Water has unique and unusual properties both in pure form and as a solvent. These properties influence chemical reactions. Structure of the water molecule: The structure of H2O is shown in Fig 3-1. It consists of an O atom with 6 e- that have the electronic configuration of: 2 2 2 1 s 2s 2pz 2py 2px which merge with two H atoms with 1 e- each resulting in a neutral molecule with 8 e- which form four pairs of electron orbitals called sp3 hybrids. The most stable configuration of these four lobes is a tetrahedral arrangement, with two e- in each lobe. Two lobes are used for O-H bonds (shared electrons) and two lobes have free lone pairs of electrons Figure 3-1 Electronic Orbitals (Fig 3-1). The water molecule lacks symmetry, the of the water molecule charge is not evenly distributed (it is polar). Electrons are shared unequally between O (highly electronegative) and H (not so greedy). The H-O-H tetrahedral angle is 105° which is less than the ideal tetrahedral angle of 109° (Fig 3-2). The reason this is so is because of e- repulsion by the lone pairs. As a result of this bent structure water has a separation of charge or a dipole moment. Thus H2O is a polar molecule. Fig 3-2 Bent Structure Fig 3-3 Hydrogen Bonding 1 As a result of this dipole moment, the water molecule has a propensity to form hydrogen bonds (H-bonds) (Fig 3-3). H-bonds often occur in liquids made up of molecules in which H is bonded to a highly electronegative element (e.g. O, N, Cl or F). 1. H-bonds can be thought of as resulting from "resonance" of an H between two more polar atoms and necessitates that the three atoms be collinear (see Fig 3-3). 2. The H-O----H bond is relatively strong and has a strength of about 4.5 kcal mol-1, compared to 10-20 kcal mol-1 for ionic bonds and ~0.5 kcal mol-1 for van der Waals bonds. These H-bonds also lead to "cooperative bonding" in which water molecules link together to actually form regions with structure. The formation of one H-bond makes it easier to form a stronger second H-bond because of the first. In some ways these regions may have an ice like structure. One line of evidence is that when ice melts only about 15% of the bonds are broken. In one model these regions are called "flickering clusters" because they form and reform at a rapid rate (the Frank-Wen Flickering Cluster Model - Fig 3-4). There are many models for the structure of water but it is not necessary to review all of them here (see Horne, 1969; Eisenberg and Kauzmann, 1968) Fig 3-4 Frank-Wen Flickering Cluster Model The structure of the H2O molecule explains why it is easier to freeze water than convert it to a gas. It also explains why H2O is a good solvent, because it is attracted to anions and cations. When a salt dissolves in solution, water molecules surround each ion in a process called hydration (see Fig 3-5). Such hydration numbers are difficult to measure but the strength of hydration is probably proportional to the charge to radius ratio (q/r) of the central ion. Thus divalent Ca2+ hydrates more strongly than monovalent Na+ and small Mg2+ hydrates more strongly than large Ca2+ . Hydration isolates ions from each other and enhances solubilization relative to other solvents (Fig 3-6). The water molecules around a central ion are called the hydration sphere. The water molecules probably arrange themselves into an inner sphere of tightly bound waters and an outer sphere that is less tightly bound. The water molecules are oriented with their oxygen atoms (or negative side) pointed toward a cation. The process of hydration usually results in a decrease in volume and is called electrostriction and 3 influences the molal volumes occupied by ions (e.g. Vion in cm /mol). 2 Example 3-1: Electrostriction and Volume The effect of hydration and electrostriction can be illustrated using a simple example. If we make a 0.5m NaCl solution we can predict the volume of the solution from the weights and densities of the recipe and compare the predicted volume with that measured. Thus: Component density volume 29.22 g NaCl ÷ 2.165 g/cm3 = 13.50 cm3 3 3 970.78g H2O ÷ 0.997 g/cm = 973.70 cm 1000.0g of 0.5m of NaCl = 987.2 cm3 (predicted) 983.0 cm3 (actual) Volume difference = 4.2 cm3 3 Fig. 3-5 Electrostriction Fig. 3-6 Hydration assists dissolution of salt crystals Fig 3-7 Boiling Point and Freezing Point for hydrides of elements (O, S, Se, Te) of group VIA in the Periodic Table 4 The anomalous properties of water Water has many unusual properties that indicate strong intermolecular associations. The structure of the H2O molecule and the hydrogen bonding of water explain many of these unique physical properties. These properties (which were originally presented by Sverdrup, Johnson and Fleming in 1942) include: high heat capacity Heat Capacity (Cp) is the thermal energy it takes to raise 1 gm of a substance by 1°C. Water has the highest heat capacity of all solids and liquids except liquid NH3. This is because it takes a lot of energy to break the hydrogen bonds and change the structure of water. Thus water has a large thermal buffer capacity and acts as a climate buffer. Energy transport by water transport in currents is large. high heat of evaporation Mammals cool by sweating! - it takes energy (∆H = 540 cal g-1) to break the hydrogen bonds. Water has the highest heat of evaporation of all liquids. high boiling point The boiling points and freezing points of the group VIA hydrides (S, Se and Te) fall on a line of decreasing B.P./F.P. with decreasing molecular weight except water (see Fig. 3-7). The projected B.P. of H2O should be - 68°C while the real value is 100°C. high freezing point It is easier to freeze water than convert it to a gas. The freezing point of 0°C is anomously high. The projected F.P. , based on the other elements of its group in the Periodic Table, is -90°C (see Fig 3-7). Water exists in 3 phases within the critical temperature range that accommodates life. low heat of freezing The water structure can move easily to the ice structure. The heat of freezing is only 1/7 that of evaporation implying that there is a relative small difference in the number of bonds between water and ice. high surface tension Water likes itself relative to most other surfaces and because of this water tends to minimize its surface area. 5 When air bubbles break at the sea surface the high surface tension causes the surrounding water to snap back into the depression left by the bubble resulting in injection of a small droplet of surface seawater into the atmosphere (Fig. 3-8). The water soon evaporates leaving a small aerosol of seasalt. This is the mechanism by which seasalt is transfered from the ocean to land. Fig. 3-8 Aerosol formation after an air bubble comes to the sea surface high dielectric constant Water has the highest dielectric constant of all substances except H2O2 and HCN. Water has a high dissolving power because the water molecules reduce the forces of attraction between ions. You could consider this as a result of ion hydration. The force between ions (F) is coulombic and the dielectric constant (ε) reduces this force according to: F = 2 q1q2/r ε , where q1 and q2 are the charges on two ions separated by a distance, r. For water at 25°C, ε = 78. The high heat capacity and heats of fusion and evaporation provide immense thermostating capacity in the critical temperature range that accomodates most life (-50 to 100 °C). a) It takes 766 calories to raise the temperature of 1 gram of water from -50 °C (as ice) to +150 °C (as steam). The same number of calories would elevate the temperature of 10 grams of granite (heat capacity = 0.2 cal g-1) to 383 °C. b) Water is unusual for having two phase transitions so close together in temperature, thereby allowing large amounts of energy to Fig. 3-9 Temperature and phases of be exchanged while the temperature of the water versus energy (as calaries) two-phase system (ice/water or water/steam) remains constant (Fig 3-9) Liquid water is best considered as a mixture of densely packed "monomers" and ice-like "polymers". Their proportions vary with temperature. As temperature decreases monomers are converted to polymers, the fraction as structured regions increases and the 6 density decreases (Fig 3-10). The maximum density for pure water (with no ions) occurs at 4°C. Ice is much less dense than water at all temperatures. Have you wondered why lakes don't freeze solid all the way to the bottom? The key is that the density of freshwater reaches a maximum at 4°C (see Fig 3-10).
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