Natural Sciences I Lecture 22: Water and Solutions

Transcription

1 1 Introduction Natural Sciences I Lecture 22: Water and Solutions Solutions figure prominently in everyday life: many of the beverages we drink, medicines we take, and household chemicals we use are actually solutions. In fact unless you buy distilled, dionized water for drinking purposes even the "pure" water you drink is really a dilute solution. For these reasons it is important to understand the nature of solutions and how we describe their properties. A solution is a homogeneous mixture of ions or molecules of two or more substances. The most abundant substance is typically referred to as the solvent; the less abundant substance(s) the solute(s). Contrary to common perception, a solu- solution type* solvent tion need not salt in water (seawater, etc.) s in l water involve a liquid CO 2 in water (soda water) g in l water at all (i.e., the solvent can be sugar in water s in l water a gas or solid). iodine in alcohol (tincture of iodine) s in l alcohol Examples: alcohol in water ("liquor") l in l water oxygen in nitrogen (air) mercury in silver (dental amalgam) zinc in copper (brass) g in g s in s s in s * s = solid; l = liquid; g = gas (room T assumed) nitrogen silver copper Solid Solutions are conceptually the trickiest. Suppose we take a crystal of NaCl and replace some of the Na ions by Li ions we have a solid solution! Cl Na Li pure NaCl NaCl with LiCl in solid solution

2 2 CONCENTRATION UNITS We often speak of solutions as being "concentrated" or "dilute", but these are qualitative indicators at best and can mean different things in different situations. Other units are needed; unfortunately there's no standardization. Solution strength by parts... parts per hundred (rarely if ever actually used) 2 parts solute 98 parts solvent 2 parts per hundred solution With reference to the diagram above, note that 2 atoms in a 10 x 10 x 10-atom cube would be 2 parts per thousand; 2 atoms in a 100 x 100 x 100-atom cube would be 2 parts per million. Parts per million (ppm) and parts per billion (ppb) are commonly used units of concentration. Unfortunately, however, they are used in different ways, sometimes referring to parts by atom (or by mole as above), sometimes to parts by mass, and occasionally even to parts by volume. The illustration at the right shows 2 parts per hundred by weight (mass) or 20,000 ppmw. The difference between this and ppm by atom can be significant if the solute atoms are very different in mass from the solvent atoms e.g., Pb in H2O

3 3 Parts by volume is sometimes used to describe solutions of one liquid in another due to ease the of volumetric measurements in the case of liquids. Seawater: parts per thousand by mass... = 965 g water 35 g salts 1000 g seawater The salinity of seawater is expressed as 35 %0 parts per thousand Other units... Some chemists use molarity (M) = Physical chemists and geochemists tend to use molality (m) = moles of solute liter of solution moles of solute kilogram of water Per cent concentrations Note that concentrations expressed in terms of atoms, mass, or volume are often converted into percentages. For example, the volume % of a liquid solute in a solution (e.g., alcohol in water) would be V solute 100% = volume % solute V solution Similarly, for a mass percentage m solute m solution 100% = mass % solute

4 4 SOLUBILITY In many types of solutions, the solute and solvent appear to be mutually soluble in all proportions. This is true of many liquids (e.g., alcohol and water) and all gases at near-atmospheric pressures (e.g., nitrogen and oxygen). There are, of course, many examples of liquids that do not mix in all proportions (e.g., oil and water). Most solids have definite limits to how much they will dissolve in water. We speak of the limit in each case as the solubility of the solute under consideration. The SOLUBILITY of a solute is defined as the concentration that is reached in a saturated solution at a particular temperature (and pressure). A saturated solution is one in which equilibrium has been reached between the pure solute and the solution, such that the amount of solute going into solution in a given time is equal to the amount that comes out of solution. It is important to realize that the equilibrium condition is not necessarily one of no change at all, but one of no net change. Equilibrium is a dynamic condition... saturated solution of NaCl in water NaCl concentration in water is constant precipitating NaCl dissolving NaCl The equilibrium is dynamic even at the atomistic scale: sodium and chlorine atoms are constantly "hopping" on and off the surface. The hops each way balance one another. NaCl crystal Na Cl Na Cl saturated solution

5 6 Attractive forces between positive and negative "poles" of individual molecules result in hydrogen bonding. In ice, hydrogen bonding leads to an open, hexagonal structure: When ice melts, about 15% of the hydrogen bonds break. The open structure collapses somewhat (the density actually increases) and the H2O molecules begin to jostle around. Hydrogen bonds remain important in influencing the properties of liquid water. structure of ice (schematic) What will dissolve in what? In a very general sense, liquids that consist of polar molecules tend to dissolve in other liquids that consist of polar molecules. Conversely, non-polar liquids tend to be insoluble in polar liquids like water. Example: Carbon tetrachloride The chlorine atoms (green) are arranged symmetrically around the carbon atom (black) to form a tetrahedron. The resulting molecule is non-polar there is no localization of charge. Cl C CCl 4 "Like dissolves like": Carbon tetrachloride is an effective solvent for other non-polar compounds such as oils and greases that are insoluble in water. Some larger molecules such as ethanol and soaps have both polar and a non-polar endts which is why they can be used with water to dissolve oily compounds.

6 5 For solids dissolving in liquids, the solubilities are highly dependent on the nature of the solid. The temperature dependencies of the solubilities also vary markedly... solubility (g solute/100 g water) solubilities of some salts in H O 2 NaNO 3 KNO 3 KCl NaCl PROPERTIES of WATER temperature (C) Most of the solutions that concern us in everyday life involve solutes in water. Accordingly, we should develop some background about the structure and properties of water in order to understand why various solutes behave they do. The first thing to recognize is that the water molecule although neutral overall has an uneven distribution of charge: We've already seen two representations... H O H and O H H o 105 Because of the high electronegativity of oxygen, the shared electrons spend more time near the oxygen atom, so the water molecule is POLAR The polarity of the water molecule leads to intermolecular forces between like molecules. These are relatively weak forces (the general term is van der Waals forces), but they have a significant influence on the properties of water. H ( ) O () H ( ) ()

7 7 The dissolving process: Atom-scale process at surfaces Up to this point, we've discussed solubilities and bonding in a general way, but as yet have no insight into what goes on at the scale of atoms when a substances dissolves in water. Let's consider NaCl as an example... The water molecules attach themselves to specific locations "sites" on the crystal surface according the polarity: the (-) "oxygen end" bonds to () sodium, the () hydrogen end side to chlorine. The overall process is called hydration. Na polar water molecules If the attraction of water molecules is strong enough relative to ionic bonding in the crystal, the ions in the crystal can be "pulled away" into the solution. At the atomistic scale, this is the process of dissolution. Ions pulled off the surface are now in solution; these are surrounded by polar water molecules to form a "sphere of hydration". The positively charged "hydrogen ends" of the water molecules are directed toward dissolved anions; the negatively charged "oxygen ends" are directed toward dissolved cations. The sphere of hydration is the region highlighted in yellow. Cl NaCl crystal DISSOLVED ANION DISSOLVED CATION

8 8 The chemical equations representing dissolution of an ionic crystal in water are written to make it clear that the reaction products are ions in aqueous solution: NaCl (s) Na (aq) Cl (aq) PROPERTIES of SOLUTIONS The properties of pure solvents such as water can be dramatically changed by the introduction of solutes. The nature and extent of the changes depends upon the identity of the solute(s) and the manner in which it dissolves. Dissolution of ionic substances (page 7) creates solutions we call electrolytes; as their name suggests, these conduct electricity. The charge is carried by mobile ions, not by electrons as in metallic conductors. 0 lots amps pure H O 2 0 amps lo ts Na Cl Cl Cl Na Na Water solutions of alcohol, sugars, and other covalently bonded molecules are nonelectrolytes they do not conduct electricity because the solutes are present in solution as neutral molecules. Schematic representation of a glucose (sugar) molecule and an ethanol (alcohol) molecule dissolved in water. The molecules are neutral but engage in hydrogen bonding with neighboring water molecules. glucose ethanol carbon oxygen hydrogen

9 9 Interestingly, some compounds that are covalently bonded as pure substances (and therefore nonelectrolytes) can dissolve in water to make electrolytes. An important example is hydrogen chloride (HCl). HCl is a colorless gas at room conditions, but condenses to a non- O conducting liquid at -85 C. When dissolved in water, HCl dissociates (becomes ionized) to make hydrochloric acid, which conducts electricity very well. HCl (g) H O (l) H O (aq) Cl (aq) 2 3 A proton and a chlorine ion are produced by the dissociation of the HCl molecule. The proton becomes associated with a water molecule to form the hydronium ion, H O Cl. water hydrogen chloride "collision complex" hydronium ion chloride ion ACIDS and BASES The preceding discussion of hydrochloric acid is a good lead-in to a more thorough treatment of acids (and bases) in general. Acids and bases are solutions characterized by specific chemical properties. We tend to think of them strictly as aqueous solutions, but it isn't necessary that the solute be water. We noted in the case of HCl is that dissolution in water produces negative chloride ions and positive hydronium ions. There is a tendency to simplify this and regard the product ions simply as chloride and hydrogen ions (H ). This is OK, but it is important to remember that H is a "bare" proton, which is very reactive and always exists in association with other ions.

10 10 Definitions of Acids and Bases Arrhenius Bronsted- Lowry Lewis an acid liberates hydrogen ions in water a base liberates hydroxide ions in water an acid donates protons to any other compound a base accepts protons from any other compound an acid can accept a share in a pair of electrons a base can donate a share in a pair of electrons too restrictive usually OK very general Let's try to get a feeling for what these definitions really mean. First, we need to broaden our perception of the nature of water. It's convenient to think of water as a collection of H2O molecules "slightly" bonded to one another by hydrogen bonds (this has been our model up to this point). The reality, however, is that at any given instant, some of the water molecules are ionized due to fortuitous collisions between two of them: H 2 O H 2 O "collision complex" H O 3 OH This reaction goes the other way, too hydronium ions recombine with hydroxide ions to form two water molecules: H 3 O OH H 2 O H 2 O

11 11 We represent the equilibrium this way, as a balanced reaction with double arrows to show that it proceeds in both directions 2H O H O OH 2 3 In pure H O at any given instant, there are a very small number of 2 hydronium and hydroxide ions present, and these must be equal. At room conditions, that number is moles/liter of H 3 O moles/liter of OH Because the hydronium and hydroxide ions balance one another, we refer to this as a neutral solution in terms of its acid-base properties no "excess" H O 3 or OH. As we saw in the case of HCl, however, dissolving a compound that contains hydrogen and that ionizes in solution can dramatically increase the concentration of hydronium. A solution with increased hydronium concentration is an acid solution. The ph can be as high as zero mole/liter of H O 3 = 1 mole/liter p H = 7 ( neutra l) (Because of the equilibrium shown at the top of the page, the product of the hydronium ion concentration and the hyroxide ion concentration -14 is always 10 ). A base is formed by ionization of a solute to produce hyroxide ions; e.g., Na O H O 2Na (aq) 2OH (aq) 2 2 This moves us off the neutral position, creating a base solution with more hydroxide ions than hydronium ions (meaning ph > 7). Given the reactions discussed above, it should make sense that the strength of an acid or base solution depends upon the extent to which the solute ionizes to produce hydronium or hyroxide ions (or other ions that can accept or donate a share in a pair of electrons).

12 12 Common acids and their extent of ionization STRONG name formula % ionization* perchloric hydrochloric hydrobromic hydroiodic nitric sulfuric HClO 4 HCl HBr HI HNO 3 H SO 2 4 ~100 ~100 ~100 ~100 ~100 ~60 very dangerous phosphoric H PO WEAK acetic carbonic boric CH3COOH H2CO3 H3BO * The % ionization depends upon concentration of the solution and T. Data are for 0.1 moles/liter and room T. an organic acid important in Earth's carbon cycle eyewash Because the general theme of this discussion is solubility, we shouldn't leave the topic of acids and bases without noting that the acidity of a solution can dramatically affect the solubilities of various compounds. This is because hydronium and hydroxide ions can aggressively attack many bonds. solubility (mg/liter) solubility of amorphous SiO 2 O in aqueous solutions at 25 C ph

13 13 Some final thoughts on solubility... Solution chemistry and solubility considerations are central to many environmental issues. In addition to exerting fundamental controls on the chemical composition of natural waters (oceans; lakes and streams; groundwater), they also determine the mobility and fate of organic and inorganic environmental contaminants ranging from heavy metals to PCBs. There are even some practical household applications... "Scale" in pipes CaCO 3 CO 2 H2O Ca(HCO 3) 2 limestone calcium bicarbonate (dissociated) HEAT 2 2HCO CO CO H O The bicarbonate ion is destroyed by heat. When water laden with calcium bicarbonate is heated (as in a hot-water system), carbonate ions are produced. However, calcium carbonate is quite insoluble in water, so it precipitate to form "scale", clogging pipes and ruining boilers. 2 2 Ca CO 3 CaCO3