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Unformatted text preview: F e I .iiiiiiIuI-:"=r= - - 8 A PROBLEM-SOLVING APPROACH'TO AQUATIC CHEMISTRY You will note in Appendix A that a low pH corresponds to high activity ole. Conditions of high l-l+ activity are called acidic. Thus, low pH (high H“ activity) means acidic conditions. High pH (low H+ activity) corre— sponds to basic conditions. Low pe (high e‘ activity) corresponds to reducing conditions, and high pe (high e’ activity) corresponds to oxidizing conditions. 1.4 PROPERTIES OF WATER Example 1.1: Solubilizafion of Sodium Chloride in Water Why does NaCl dissolve in water? Solution: One reason NaCl dissolves in water is that the potential en- ergy of interaction between Na+ and Cl‘ is reduced in water. To illustrate this point, calculate the potential energy of interaction between Na+ and Cl" separated by 1 nm in water. The potential energy of two particles of charge ql and q 2 separate by a distance r is given by Coulomb’s Law: potential energy =q1q/(4nreae) where EU is the permittivity of a vacuum (= 8.854X10'12 J‘lsz'l) and e is the dielectric constant (= i for a vacuum and 80 for water). The unit charge on Na+ is +1.602x10'19 c and —1.602><10'19 C on Cl‘. 1.4.1 Introduction r Water is a unique substance. In freshman chemistry courses, you probably learned of the periodic nature of the properties of the elements. Elements in the same column of the periodic table share similar chemical properties. Yet water, H20, has very different physical and chemical properties than other dihydrogen complexes of elements in oxygen’s column of the periodic table. For example, water is a liquid under standard temperature and pressure conditions, but H28 is a gas. The unique properties of water stem from the large differences between the affinity of oxygen and hydrogen for electrons. Oxygen has a much higher affinity for electrons than hydrogen, resulting in a highly polarized bond between 0 and H. In fact, the oxygen atoms in water are partially negatively charged and the hydrogen atoms are partially positively charged, _ as illustrated in Figure 1.3. As a result of this polar bond and the difference in partial charge, oxygen atoms in water have the ability to form weak, but important, chemical bonds with more that two hydrogen atoms. These weak bonds are called hydrogen bonds. Without exaggeration, hydrogen bonds influence every molecular interaction in the aquatic environmentl Hydrogen bonds affect the three-dimensional shape of water. The bond distance between oxygen in water and the hydrogen-bonded hydrogen is about twice the bond distance between the oxygen in water and one of its own hydrogen atoms. Hydrogen bonding allows for the formation of very large clusters of water molecules (about 400 molecules per cluster; Luck, 1998). The clusters are shifting constantly since the average lifetime of a hydrogen bond is only a few picoseconds (10’12 s). The shifting clusters of water molecules contribute to the solubilization of ions in water, as will be discussed in Section 1.4.2. T Consider, for example, the influence of hydrogen bonds on ice. The structure of ice is different than the structure of water. In ice, all the water molecules participate in four hydrogen bonds arranged in a tetrahedronal shape. The shape results in an open structure for ice and a corresponding density of 0.92 g/cm3. The density of ice is less than the density of liquid water at 0°C, so ice floats. The length of the hydrogen bond is directly responsible for the open structure and low density of ice. If the hydrogen bond was 5% shorter, ice would sink and water bodies would have frozen solid during the ice ages. If hydrogen bonds were slightly shorter, life on Earth may not have survived long enough to allow you to read this book; — l. GETTIN Plugging in the values, you will find that the potential energy is 5_ -139 1d per mole in a vacuum 0 and —1.7 kl per mole in water /<>\ 1 04.50 l (the negative sign indicates an energy of attraction and one mole is 6.022x1023 ions). Figure 1.3: Charge Distribution Among the Atoms in Water The attractive potential energy of interaction between Na+ and Cl" in water is very small (only about one-tenth of the energy of I | l a hydrogen bond). 1.4.2 Solubility of ionic species I Throughout this text, it shall be assumed that saits containing sodium, ‘ potassium, or chloride at low concentrations dissolVe nearly completely in water to form ions. Why are salts so soluble in water? Hydrogen bonding allows for the orientation of a large number of molecules simultaneously. 1 i This orientation reduces the field strength of an applied field significantly. ' More specifically, the orientation reduces the attractive forces betwoen pairs of anions and cations. As the attractive forces are reduced, the salt can be so‘lubilized. 1For example, you can show that water significantly reduces the attractive forces between Na+ and Cl‘ (sec Example 1.1). Sodium chloride is soluble in water mainly because the attractive energy between Na+ and Cl” is reduced by water. (Note: The energy of the system also is altered as the atoms in the salt become more disordered and the water molecules be— come more ordered around the 1.4.3 Hydration hydrated. In other words, the ions form Once dissolvod, the ions become molecules (called complexes) with water. Water is adept at hydrating both cations and anions. Why? Recall from Section 1.41 that water has both a partially positive pole on H and partially negativo pole on 0. Water is said to be very polar. For example, when alum is added to water at low pH, AP+ forms and is quickly hydrated to Al(H20)53+. Such structures are frequently For example, sometimes Cu2+ is written to represent ions.) ______.__—————— hydration: the process of form- ing compounds with water polar: a polar species contains atoms or groups exhibiting abbmviated. different charges (or partial Cu(H20)42+l Charges) Even the proton is hydrated. Free I-I+ probably does not exist in solution. The free proton is hydrated to form H(HZO)+ = H30“, P101120); : H50; and others. For convenience, the collection of hydrated protons is abbreviated as H+ or H30“. ' 1.5 PART I ROAD MAP Part I of this book contains the fundamental concepts that must be mastered before calculating the equilibrium concentrations of chemical species. Readers with a chemistry background limited to freshman chemistry only should study each Section of each chapter and work the problems assigned. In Chapter 2, concentration units are reviewed. The main take-home lesson from Chapter 2 is the definition of molar concentration units and why we use molar concentration units for equilibrium calculations. ...
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This note was uploaded on 02/02/2012 for the course CEE 254 taught by Professor Jenniferjay during the Fall '11 term at UCLA.

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JensenPropofWater - F e I .iiiiiiIuI-:"=r= - - 8 A...

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