1. Ch. 18 Properties of Solutions

2. CA Standards

3. Let’s review some things before we start IonicMolecular Types of atoms that form the bond Cations – metals from groups 1-12 Anions – nonmetals Nonmetals for all atoms involved How the bond is formed Electrons are stolen by the anion (from the cation). Electrons are shared, but if it is a polar bond, shared unequally (penguin and polar bear) ExamplesNaCl, CaCO 3, ZnCl 2 CH 4, NH 3, H 2 O Naming and formulasCation uses name of atom, anion add –ide ending Use di, tri, tetra, etc. prefixes to state how many of each.

4. Review: Intermolecular forces Dispersion forces – the weakest kind of intermolecular force; this attraction is thought to be caused by the motion of electrons. Dipole interaction (dipole-dipole) – a weak intermolecular force resulting from the attraction of oppositely charged regions of polar molecules ( δ + to δ - attraction ) Hydrogen bonding – a relatively strong intermolecular force in which a H atom that is covalently bonded to an electronegative atom (meaning the H is an almost naked proton) is also weakly bonded to an unshared electron pair on a neighboring molecule.

5. Classification of Matter Solutions Solutions are homogeneous mixtures

6. 18.1 Properties of Solutions We often think of solutions as being a solid dissolved in a liquid, and that is often the case. But solutions can be composed of solid, liquid or gas. Here’s some examples: Some common types of solutions Gas – gasCarbon dioxide and oxygen in nitrogen (air) Liquid – gasWater vapor in air (moist air) Gas – liquidCarbon dioxide in water (soda water) Liquid – liquidAcetic acid in water (vinegar) Solid – liquidSodium chloride in water (salt water/brine) Solid - solidCopper in silver (sterling silver alloy) AqueousA solute that is dissolved in water

7. Solute A solute is the dissolved substance in a solution. A solvent is the dissolving medium in a solution. Solvent Salt in salt waterSugar in soda drinks Carbon dioxide in soda drinks Water in salt water Water in soda Consider a solution of solid dissolved in liquid:

8. Here the green and yellow solute atoms are an ionic solid, say NaCl. The red and white molecules are the polar solvent, water. Do you notice how the water molecules align themselves differently around the yellow cations and the green anions?

9. “Like Dissolves Like” Fats Fats Benzene Benzene Steroids Steroids Hexane Hexane Waxes Waxes Toluene Toluene Polar and ionic dissolve best in polar ionic solutes solvents Nonpolar dissolve best in nonpolar solutes solvents Inorganic Salts Water Water Sugars Sugars Small alcohols Small alcohols Acetic acid Acetic acid

10. Solubility Trends   The solubility of MOST solids increases with temperature.   The rate at which solids dissolve increases with increasing surface area of the solid.   The solubility of gases decreases with increases in temperature.   The solubility of gases increases with the pressure above the solution.

11. Therefore… Solids tend to dissolve fastest/best when: o Heated o Stirred o Ground into small particles Gases tend to dissolve best when: o The solution is cold o Pressure is high

12. Solubility Review: Solubility is a surface phenomenon. Agitation causes the solute to dissolve faster because it brings fresh solvent in contact with the surface of the solute. However, this won’t change the solubility of a substance. If a substance is insoluble, no amount of shaking or stirring will help. Higher temperature helps the solute dissolve because the molecules of solvent have more kinetic energy when they collide with the solute, so there is an increased frequency and force of collisions with the solute particles. Grinding the solute into small particles increases the surface area of the solute, allowing a faster reaction. Think of a sugar cube vs. granulated sugar.

13. Review Three things that affect HOW FAST a substance dissolves: – –Stirring/agitation – –Temperature – –Surface area Three things that affect SOLUBILITY of a substance: – –Most substances higher solubility at higher temperatures. – –The nature of the solute and solvent itself. – –For gases in liquids, higher pressure increases solubility.

14. Solubility limits Many solutes have limited solubility in a given solvent. For example, have you ever poured a whole bunch of salt into a cup of water? Did you notice there was still a lot of undissolved salt at the bottom of the cup? The solubility limit for NaCl in 25 o C water is 36.2 g per 100 g of water. If you add more than that, it will just sit at the bottom of the container as a solid and not go into solution. An unsaturated solution contains solute that is LESS than the solubility limit. A saturated solution is one that contains the maximum amount of solute for a given amount of solvent at a given temperature (at the solubility limit or higher).

15. Solute above the solubility limits The solubility is the maximum amount of solute that will dissolve in a given solvent at a given temperature. But once you get above that limit, in a saturated solution, there is actually a dynamic equilibrium between molecules of solute going into solution and falling out of solution (equal rates of solvation and crystallization)

16. Solubility Chart What is the solubility of NaCl at 70 o C? At 25 o C, how many grams of potassium nitrate will dissolve in 200 mL water?

17. PhET solubility simulator soluble-salts_en.jar

18. Supersaturated solutions When the temperature of a saturated solution in contact with a small excess of solid solute is raised, then some of that excess solute dissolves because of the raised temperature. If the solution is then cooled, some of that excess solute may not immediately crystallize out of solution. That leaves you with a supersaturated solution. If you now add a seed crystal, then the excess solute may rapidly crystallize out on that seed crystal. This is how rock candy is made.

19. Solubility of gases - temperature Notice that in general the solubility of gases goes down as temperature increases, to nearly zero at 100 o C.

20. Solubility of gases: pressure In the sealed bottle, the high pressure of CO 2 above the liquid keeps the concentration of dissolved CO 2 high. But when the cap is removed, the equilibrium is disrupted and the CO 2 pressure and the solubility of CO 2 goes down.

21. Henry’s Law (liquid-gas)

22. Miscibility (liquid-liquid) What about liquid-and-liquid solutions? Liquids are said to be miscible if they dissolve in each other. – –Example of miscible liquids – water and ethanol. – –Example of immiscible liquids – oil and vinegar, because they separate from each other and will not mix. Oil and vinegar on left. Oil spill on a body of water, not mixing, on right.

23. An electrolyte is: A substance whose aqueous solution conducts an electric current. All ionic compounds are electrolytes. Sodium chloride, copper(II) sulfate, sodium hydroxide are examples. Many acids/bases too. A nonelectrolyte is: A substance whose aqueous solution does not conduct an electric current. Most molecular compounds are not composed of ions, so they are nonelectrolytes. Examples: sugar, rubbing alcohol, C-compounds Definition of Electrolytes and Nonelectrolytes

24. The ammeter measures the flow of electrons (current) through the circuit. If the ammeter measures a current, and the bulb glows, then the solution conducts. If the ammeter fails to measure a current, and the bulb does not glow, the solution is non-conducting. Electrolytes vs. Nonelectrolytes

25. 1.Pure water 2.Tap water 3.Sugar solution 4.Sodium chloride solution 5.Hydrochloric acid solution 6.Lactic acid solution 7.Ethyl alcohol solution 8.Pure sodium chloride 1.Pure water 2.Tap water 3.Sugar solution 4.Sodium chloride solution 5.Hydrochloric acid solution 6.Lactic acid solution 7.Ethyl alcohol solution 8.Pure sodium chloride Try to classify the following substances as electrolytes or nonelectrolytes…

26. ELECTROLYTES: NONELECTROLYTES: Tap water (weak) NaCl solution HCl solution Lactate solution (weak) Pure water Sugar solution Ethanol solution Pure NaCl Answers to Electrolytes

27. 18.2 Concentrations of Solutions The concentration of a solution is the measure of the amount of solute that is dissolved in a given amount of solvent.

28. Concentration A dilute solution is one that contains only a low concentration of solute. Ex: 1 g NaCl per 100 g H 2 O. A concentrated solution is one that contains a high concentration of solute. Ex: 30 g NaCl per 100 g H 2 O. Molarity (M) is the number of moles of a solute divided by liters of solution. Molarity is also known as molar concentration. Example: I have a 0.5 molar solution of NaCl means my solution has 0.5 moles of NaCl per liter of H 2 O (assuming the solvent is water).

29. Making a 0.5 molar solution of NaCl in water a.Add 0.5 mol solute to a 1 L volumetric flask half filled with distilled water. Since the molar mass of NaCl is 58.5 g/mol, you would measure 58.5/2 = 29.25 g of NaCl. b.Swirl the flask carefully to dissolve the solid. c.Then fill the flask to the 1 L mark. Now you have a 0.5 molar solution.

30. Examples with molarity

31. Example 18-2

32. Example 18-3

33. How do you make a dilution? If you want to make a dilution, that means you are going to take a more concentrated solution and dilute it such that there are fewer molecules of solute per liter of solvent, but the overall number of molecules of solute doesn’t change.

34. Dilutions

35. PhET concentration simulator concentration_en.jar

36. Example 18-4

37. Now you try one:

38. How about this one:

39. Percent Solutions (liquid-liquid)

40. Percent Solutions (solid-liquid)

41. Let’s try Example 18-5

42. Example 18-6

43. Now you try it

44. Parts per million / billion