Chapter Nine

1. Fundamentals of General, Organic and Biological Chemistry 5th Edition Chapter Nine Solutions James E. Mayhugh Oklahoma City University 2007 Prentice Hall, Inc.

2. Outline • 9.1 Mixtures and Solutions • 9.2 The Solution Process • 9.3 Solid Hydrates • 9.4 Solubility • 9.5 The Effect of Temperature on Solubility • 9.6 The Effect of Pressure on Solubility: Henry’s Law • 9.7 Units of Concentration • 9.8 Dilution • 9.9 Ions in Solution: Electrolytes • 9.10 Electrolytes in Body Fluids: Equivalents and Milliequivalents • 9.11 Properties of Solutions • 9.12 Osmosis and Osmotic Pressure • 9.13 Dialysis Chapter Nine

3. 9.1 Mixtures and Solutions • Heterogeneous mixture: A nonuniform mixture that has regions of different composition. • Homogeneous mixture: A uniform mixture that has the same composition throughout. • Solution A homogeneous mixture that contains particles the size of a typical ion or small molecule. • Colloid A homogeneous mixture that contains particles in the range 2–500 nm diameter. Chapter Nine

4. Liquid solutions, colloids, and heterogeneous mixtures can be distinguished in several ways. Chapter Nine

5. 9.2 The Solution Process A good rule of thumb for predicting solubility is that “like dissolves like”. Substances with similar intermolecular forces form solutions and substances with different intermolecular forces do not. Chapter Nine

6. Dissolution of an NaCl crystal in water. Polar water molecules surround individual ions pulling them from the crystal surface into solution. Oxygen atoms point to (+) ions and hydrogen atoms point to (-) ions. Chapter Nine

7. 9.3 Solid Hydrates • Some ionic compounds attract water strongly enough to hold onto water molecules even when crystalline, forming what are called solid hydrates. • Plaster of Paris, CaSO4·1/2H2O, is a solid hydrate. The formula indicates that for every 2 formula units of calcium sulfate in the crystal there is also one water. • Still other ionic compounds attract water so strongly that they pull water vapor from humid air to become hydrated. Compounds that show this behavior, such as calcium chloride are called hygroscopicand are often used as drying agents. Chapter Nine

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9. 9.4 Solubility • Hydrogen bonding between water and ethanol, between water alone, and ethanol alone is so similar that the two liquids are miscible, or mutually soluble in all proportions. • Most substances reach the limit of a saturated solution: A solution that contains the maximum amount of dissolved solute at equilibrium. • Solubility: The maximum amount of a substance that will dissolve in a given amount of solvent at a specified temperature. Chapter Nine

10. 9.5 The Effect of Temperature on Solubility • Temperature often has a dramatic effect on solubility. • The effect of temperature is different for every substance, however, and is usually unpredictable. • Solids that are more soluble at high temperature than at low temperature can sometimes form what are called supersaturated solutions, which contain even more solute than a saturated solution. • Such a solution is unstable and precipitation can occur dramatically when a tiny seed crystal is added. Chapter Nine

11. Solubility of some (a) solids and (b) gases, in water as a function of temperature. • Most solid substances become more soluble as temperature rises. • The solubility of gases decreases as temperature rises. Chapter Nine

12. 9.7 Units of Concentration • Solute: A substance dissolved in a liquid. • Solvent: The liquid in which a substance is dissolved. • Solution: The combination of solute and solvent. • A very useful means of expressing concentration in the laboratory is molarity (M), the number of moles of solute dissolved per liter of solution. Chapter Nine

13. Problem 9.8 What is the molarity of a solution that contains 50.0 g of vitamin B1 hydrochloride (molar mass = 337 g) in 160 mL of solution? Chapter Nine

14. Problem 9.8 What is the molarity of a solution that contains 50.0 g of vitamin B1 hydrochloride (molar mass = 337 g) in 160 mL of solution? Molarity = Moles solute/Liters solution Moles of solute = 50.0 g x 1 mol/337 g = 0.148 mol Liters of solution = 160 mL x 1 L/1000 mL =0.160 L Molarity = 0.148 mol/0.160 L = 0.925 M Chapter Nine

15. Problem 9.9 How many moles of solute are in 175 mL of 0.35M NaNO3? Volume (L) xMolarity (moles/L) = moles Chapter Nine

16. Problem 9.9 How many moles of solute are in 175 mL of 0.35M NaNO3? Volume (L) xMolarity (moles/L) = moles Volume = 175 mL x 1 L/1000 mL = 0.175 L Moles =0.175 L x 0.35mol/L = 0.061 mol NaNO3 Chapter Nine

17. Problem 9.10 The concentration of cholesterol in blood is approximately 5.0mM. How many moles of cholesterol are in 1.0 L of blood? How many grams? Chapter Nine

18. Problem 9.10 The concentration of cholesterol (C27H46O) in blood is approximately 5.0mM. How many moles of cholesterol are in 1.0 L of blood? How many grams? 5.0mM = 5.0 mmol/L x 1 mol/1000 mmol = 0.0050 mol/L (0.0050M) How many grams? 0.0050 mol x 386.7 g / 1 mol = 1.93 g cholesterol in 1 L of blood Chapter Nine

19. Cholesterol Levels Cholesterol levels typically measured in mg cholesterol / dL blood 1.93 g x 1000 mg / 1 g = 1.93 x 103 mg 1 L x 10 dL / 1L = 10 dL 1.93 x 103 mg= 193 mg/dL 10 dL Chapter Nine

20. Weight/Volume Percent Concentration [(w/v)%] • Mathematically, (w/v)% concentration is found by taking the number of grams of solute per milliliters of solution and multiplying by 100. • Physiological saline solution: 0.90% NaCl 0.90 g NaCl dissolved and diluted to a total volume of 100 mL Chapter Nine

21. Dissolving one liquid in another Volume/Volume Percent Concentration [(v/v)%] Mathematically (v/v)% is determined from the volume of solute (usually in mL) per mL of solution multiplied by 100. What is the v/v% of a 75.0 mL solution made with 3.80 mL methanol? v/v% = Vsolute/Vsolution x 100 = 3.80 mL/75.0 mL x 100 = 5.07% Chapter Nine

22. Parts per Million (ppm) and Parts per Billion (ppb): When concentrations are very small, as often occurs in dealing with trace amount of pollutants or contaminants, parts per million (ppm) or parts per billion (ppb) units are used. Chapter Nine

23. To prepare 100 mL of a specific solution, the solute is measured out and dissolved in just enough solvent to give a final volume of 100 mL. • If the solute were dissolved in 100 mL of solvent, the final volume of the solution will likely be a bit larger or smaller than 100 mL. Chapter Nine

24. 9.8 Dilution • Dilution: Lowering concentration by adding additional solvent. • Dilution factor: The ratio of the initial and final solution volumes (V1/V2). • In the dilution process, the amount of solute remains constant, only the volume is increased. • Moles of solute = M1V1 = M2V2 = constant • Dilution equations can be generalized to other concentration units, C1V1 = C2V2 Chapter Nine

25. The following equation is very useful in calculating final concentration of a solution after dilution. • M1 V1 =M2V2 • M1 and V1 refers to the initial concentration and volume of the solution and M2 and V2 refers to the final concentration and volume of the solution. • The final concentration will be equal to the product of the initial concentration and the dilution factor. • M2 = M1 · (V1/V2). Chapter Nine

26. Problem 9.19 Hydrochloric Acid (HCl) is normally purchased at a concentration of 12.0M. What is the final concentration if 100 mL of 12.0M HCl is diluted to 500 mL? C1V1 = C2V2(solve for C2) C2 = C1V1/V2 = (12.0M)(100mL)/500 mL = 2.40M HCl Chapter Nine

27. 9.9 Ions in Solution: Electrolytes • Electrolyte: A substance that produces ions and therefore conducts electricity when dissolved in water. • Strong electrolyte: A substance that ionizes completely when dissolved in water. • Weak electrolyte: A substance that is only partly ionized in water. • Nonelectrolyte: A substance that does not produce ions when dissolved in water. Chapter Nine

28. 9.10 Electrolytes in Body Fluids: Equivalents and Milliequivalents • What happens if NaCl and KBr are dissolved in the same solution? The cations and anions are all mixed together so an identical solution could just as well be made from KCl and NaBr. We can only speak of having a solution with four different ions in it. • A similar situation exists for blood and other body fluids, which contain many different anions and cations. Since they are all mixed together, it is difficult to talk about specific ionic compounds. • Instead, we are interested only in individual ions and in the total numbers of positive and negative charges. We need a new term, equivalentsof ions. Chapter Nine

29. One equivalent (Eq) of an ion is an amount equal to the molar mass of the ion divided by the number of its charges: 1 milliequivalent (mEq) = 0.001 equivalent (Eq) 1 Eq = 1000 mEq Chapter Nine

30. 9.11 Properties of Solutions • The properties of solutions are similar in many respects to those of pure solvents, but there are some important differences. • The properties of a solution that depend on the concentration of a dissolved solute but not on its identity are known as colligative properties. • Vapor pressures are lower, boiling points are higher and freezing points are lower in solutions than they are in a pure solvent. Chapter Nine

31. For each mole of solute particles added the boiling point of 1 kg of water is raised by 0.51°C. 1 mol of a molecular substance like glucose raises the boiling point of 1 kg of water from 100.0°C to 100.51°C. The addition of 1 mol of NaCl per kilogram of water raises the boiling point by 2 * 0.51°C = 1.02°C because the solution contains 2 mol of solute particles. • For each mole of solute particles, the freezing point of 1 kg of water is lowered by 1.86°C. 1 mol of antifreeze per kilogram of water lowers the freezing point from 0.00°C to - 1.86°C and 1 mol of NaCl (2 mol of particles) per kilogram lowers the freezing point from 0.00° C to -3.72°C. Chapter Nine

32. A close-up plot of vapor pressure versus temperature for pure water (red curve) and for a 1.0 M NaCl solution (green curve). Pure water boils at 100.0°C, but the solution does not boil until 101.0°C. Chapter Nine

33. A mixture of salt and ice is used to provide the low temperatures needed to make old-fashioned hand-cranked ice cream. • Daniel Fahrenheit did not choose 32 degrees as the freezing point of water, he chose 0 degrees to be the lowest temperature he could reach in his lab by mixing ice, water and salt. Chapter Nine

34. 9.12 Osmosis and Osmotic Pressure • Osmosis: The passage of solvent through a semipermeable membrane separating two solutions of different concentration. • Osmotic pressure: The amount of external pressure applied to the more concentrated solution to halt the passage of solvent molecules across a semipermeable membrane. Chapter Nine

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36. Osmolarity (osmol): The sum of the molarities of all dissolved particles in a solution. • Isotonic: Having the same osmolarity. • Hypotonic: Having an osmolarity less thanthe surrounding blood plasma or cells. • Hypertonic Having an osmolarity greater thanthe surrounding blood plasma or cells. Chapter Nine

37. In (a) an isotonic solution, 0.30 osmol, the blood cells are normal in appearance. • The cells in (b) a hypotonic solution are swollen because of water gain, and may burst, a process called hemolysis. • Those in (c) a hypertonic solution are shriveled because of water loss, this process is called crenation. Chapter Nine

38. 9.13 Dialysis • Dialysisis similar to osmosis, except that the pores in a dialysis membrane are larger than those in an osmotic membrane so that both solvent molecules and small solute particles can pass through, but large colloidal particles such as proteins cannot pass. • Hemodialysisis used to cleanse the blood of patients whose kidneys malfunction. Blood is diverted from the body and pumped through a long cellophane dialysis tube suspended in an isotonic solution formulated to contain many of the same components as blood plasma. Chapter Nine

39. Small waste materials such as urea pass through the dialysis membrane from the blood to the solution side where they are washed away. Chapter Nine

40. Chapter Summary • Mixtures are classified as either heterogeneous or homogeneous. Solutions are homogeneous mixtures that contain particles the size of ions and molecules whereas larger particles (2.0–500 nm diameter) are present in colloids. • The maximum amount of a solute that can be dissolved in a solvent is called the solubility. Substances tend to be mutually soluble when their intermolecular forces are similar. • The solubility in water of a solid often increases with temperature, but the solubility of a gas decreases with temperature. Chapter Nine

41. Chapter Summary Contd. • Pressure significantly affects gas solubility, which is directly proportional to the partial pressure over the solution (Henry’s law). • The concentration of a solution can be expressed in several ways, including molarity, weight/weight percent composition, weight/volume percent composition, and parts per million. • Molarity, the number of moles of solute per liter of solution, is the most useful method when calculating quantities of reactants or products for reactions in aqueous solution. Chapter Nine

42. Chapter Summary Contd. • A dilution is carried out by adding more solvent to an existing solution. Only the amount of solvent changes; the amount of solute remains the same. • Substances that form ions when dissolved in water and whose water solutions therefore conduct an electric current are called electrolytes. • Strong electrolytes ionize completely in water, weak electrolytes ionize partially, and nonelectrolytes do not ionize in water. • Fluids containing many different electrolytes have concentrations expressed in equivalents. Chapter Nine

43. Chapter Summary Contd. • A solution has a lower vapor pressure, a higher boiling point, and a lower melting point than a pure solvent. • Colligative properties depend only on the number of dissolved particles, not on their chemical identity. • Osmosisoccurs when solutions of different concentration are separated by a semipermeable membrane that allows solvent molecules to pass but blocks the passage of solute ions and molecules. • In dialysis, the membrane allows the passage of solvent and small dissolved molecules but prevents passage of proteins and larger particles. Chapter Nine

44. END OF CHAPTER 9 Chapter Nine