18.1 PROPERTIES OF SOLUTIONS

1. Chapter 18: Solutions • 18.1 PROPERTIES OF SOLUTIONS • Solution Formation – Recall that Solutions are homogenous (physical) mixtures • Solute dissolved particles in a solution (lesser amount of the two parts) are dissolved in the Solvent (the greater amount of the two parts) forming a solution.

2. How well a material (solute) dissolves in a solvent is described by a substance’s Solubility. • There are three factors that determine how fast a substance dissolves are: stirring (agitation), temperature and surface area. • Think about ways you have used these factors to better dissolve your sugar into your tea.

3. Temperature effects Solubility SOLIDS

4. Temperature effects Solubility GASES

5. Salt Solubility Simulation by PhET & Colorado University

6. Water: The Super Solvent • Water is an excellent solvent. It is used to dissolve most of the substances you have encountered in your life. In fact it is the solubility of the water in your body, blood and cells that accounts for most of the chemical processes that allow you to stay alive. • Most of the water on Earth is not pure, but rather is present in solutions.

7. Water: The Super Solvent • Water is such a versatile solvent that it is sometimes called the universal solvent. • Water is difficult to keep pure because it is an excellent solvent for a variety of solutes. • Its ability to act as a solvent is one of its most important physical properties. • As you will see, it is again the attraction of water molecules for other molecules, as well as for one another, that accounts for these solvent properties.

8. Water Dissolves Many Ionic Substances • Salt, like a great many ionic compounds, is soluble in water. The salt solution is also an excellent conductor of electricity. • This high level of electrical conductivity is always observed when ionic compounds dissolve to a significant extent in water. • Your model of water and its interactions explains why salt and many other ionic compounds dissolve in water and why the solutions conduct electricity.

9. A Model of the Dissolving of NaCl • Remember that ionic solids are composed of a three-dimensional network of positive and negative ions, which form strong ionic bonds.

10. A Model of the Dissolving of NaCl • The process by which the charged particles in an ionic solid separate from one another is called dissociation. Click box to view movie clip.

11. A Model of the Dissolving of NaCl • You can represent the process of dissolving and dissociation in shorthand fashion by the following equation.

12. Water Dissolves Many Covalent Substances • Water is not only good at dissolving ionic substances. It also is a good solvent for many covalent compounds. • Consider the covalent substance sucrose, commonly known as table sugar, as an example.

13. Water Dissolves Many Covalent Substances • You have probably observed that this substance, with the formula C12H22O11, dissolves in water. In fact, it is highly soluble. • It’s possible to dissolve almost 200 g of sugar in 100 mL of water.

14. Water Dissolves Many Covalent Substances • Take a look at the molecular structure of a sucrose molecule. • Notice that the structure has a number of O—H bonds.

15. As you learned earlier, if a molecule contains O—H bonds, it will tend to be polar and it can form hydrogen bonds. Water Dissolves Many Covalent Substances

16. Water: The Molecular View • Because water is so much a part of life, its properties are easy to take for granted. • If you step back a bit and examine water scientifically, you will find that it is unusual among the compounds found on Earth. • Water is most often thought of as a liquid.

17. Water: The Molecular View • However, solid water, called ice, and gaseous water, called steam or water vapor, also exist in large quantities on Earth. • Water is the only substance on Earth that exists in large quantities in all three states.

18. Geometry of the Water Molecule • The arrangement of electrons about the central oxygen in the water molecule relates to its three-dimensional geometry. • There is a large electronegativity difference between the covalently bonded hydrogen and oxygen.

19. Because of the molecule’s bent structure, the poles of positive and negative charge in the two bonds do not cancel, and the water molecule as a whole is polar. Geometry of the Water Molecule • Therefore, the electron pair is shared unequally.

20. Like Dissolves Like • Although water dissolves an enormous variety of substances, both ionic and covalent, it does not dissolve everything. • The phrase that scientists often use when predicting solubility is “like dissolves like.” • The expression means that dissolving occurs when similarities exist between the solvent and the solute.

21. Concentrated Versus Dilute • Chemists never apply the terms strong and weak to solution concentrations. • As you’ll see in the next chapter, these terms are used in chemistry to describe the chemical behavior of acids and bases. • Instead, use the terms concentrated and dilute.

22. Unsaturated Versus Saturated • Another way of providing information about solution composition is to express how much solute is present relative to the maximum amount the solution could hold. • If the amount of solute dissolved is less than the maximum that could be dissolved, the solution is called an unsaturated solution.

23. Unsaturated Versus Saturated • Such a solution, which holds the maximum amount of solute per amount of the solution under the given conditions, is called a saturated solution.

24. Unsaturated Versus Saturated • An interesting third category of solution is called a supersaturated solution. • Such solutions contain more solute than the usual maximum amount and are unstable. • They cannot permanently hold the excess solute in solution and may release it suddenly.

25. Unsaturated Versus Saturated • Supersaturated solutions, as you might imagine, have to be prepared carefully. • Generally, this is done by dissolving a solute in the solution at an elevated temperature, at which solubility is higher than at room temperature, and then slowly cooling the solution.

26. Effect of Temperature on Solubility • Temperature has a significant effect on solubility for most solutes. • The solubilities of some solutes, such as sodium nitrate and potassium nitrate, increase dramatically with increasing temperature.

27. Effect of Temperature on Solubility • Other solutes, like NaCl and KCl, show only slight increases in solubility with increasing temperatures. • A few solutes, like cerium(III) sulfate, Ce2(SO4)3, decrease in solubility as temperature increases.

28. Molarity • Concentration units can vary greatly. • They express a ratio that compares an amount of the solute with an amount of the solution or the solvent. • For chemistry applications, the concentration term molarity is generally the most useful. • Molarity is defined as the number of moles of solute per liter of solution. • Molarity = moles of solute/liter of solution

29. Molarity • Note that the volume is the total solution volume that results, not the volume of solvent alone. • Suppose you need 1.0 L of the salt solution mentioned above.

30. Molarity • In order to be at the same concentration as the salt in the patient’s blood, it needs to have a concentration of 0.15 moles of sodium chloride per liter of solution. • In other words, it must have a molarity of 0.15.

31. Molarity • To save space, you refer to the solution as 0.15M NaCl, where the M stands for “moles/liter” and represents the word molar. • Thus, you need 1.0 L of a 0.15-molar solution of NaCl. How are you going to prepare it?

32. Molarity • Assuming you’re making an aqueous solution, you need to know only three things when working quantitatively: the concentration, the amount of solute, and the total volume of solution needed.

33. Preparing 1 L of an NaCl Solution • How would you prepare 1.0 L of a 0.15M sodium chloride solution? • First, determine the mass of NaCl to add to a 1.0-L container. • The 0.15M solution must contain 0.15 moles of NaCl per liter of solution.

34. Preparing 1 L of an NaCl Solution • The proper setup, showing the conversion factors, is as follows.

35. Preparing 1 L of an NaCl Solution • Then carry out cancellations and calculate the answer.

36. Preparing 1 L of an NaCl Solution • The result means you need to measure 8.8 g of NaCl, add some water to dissolve it, and then add enough additional water to bring the total volume of the solution to 1.0 L.

37. Preparing a Different Volume of a Glucose Solution • How would you prepare 5.0L of a 1.5M solution of glucose, C6H12O6? • You need to determine the number of grams of glucose to add to a 5.0-L container.

38. Preparing a Different Volume of a Glucose Solution • The 1.5M solution must contain 1.5 mol of glucose per liter of solution. • The proper setup, showing the conversion factors, is as follows.

39. Preparing a Different Volume of a Glucose Solution • Cancel units and carry out the calculation.

40. Preparing a Different Volume of a Glucose Solution • The mass of glucose required is 1400 g. • Weigh this mass, add it to a 5.0-L container, add enough water to dissolve the glucose, and fill with water to the 5.0-L mark.

41. Calculating Molarity • You add 32.0 g of potassium chloride to a container and add enough water to bring the total solution volume to 955 mL. What is the molarity of this solution? • You are given that there are 32.0 g of solute per 955 mL of solution, so this relationship can be expressed in fraction form with the volume in the denominator.

42. Calculating Molarity • Therefore, the initial part of the setup is as follows.

43. Calculating Molarity • Determine that the molar mass of KCl is 74.6 g/mol by adding the atomic masses of K and Cl and applying the unit grams/mole to the sum. • The conversion factor that must be used to convert from grams to moles of KCl is 1mol KCl/74.6 g KCl.

44. Calculating Molarity • Next, to convert milliliters to liters, given that there are 1000 mL solution/L solution, use that conversion factor in the setup.

45. Calculating Molarity • Cancel units and carry out the calculation, using the setup just developed.

46. Freezing-Point Depression • A solution always has a lower freezing point than the corresponding pure solvent. • If you are interested only in aqueous solutions, this means that any aqueous solution will have a freezing point lower than 0°C. • The amount that the freezing point is depressed relative to 0°C depends only upon the concentration of the solute.

47. Boiling-Point Elevation • You have just learned that the freezing point of a solution is lower than the freezing point of the pure solvent. • It turns out that the boiling point of a solution is higher than the boiling point of the pure solvent.

48. Boiling-Point Elevation • For aqueous solutions this means that the solution boiling point will be greater than 100°C, assuming standard atmospheric pressure. • The solute must also be nonvolatile; that is not able to evaporate readily.

49. Molality • The molality (m) of a solution is equal to the number of moles of solute per kilogram of solvent.

50. Calculating Molality • What is the molality of a solution that contains 16.3 g of potassium chloride dissolved in 845 g of water? • Convert the mass of solute to moles.