Chapter 9

1. Solutions Chapter 9

2. Mixtures • A mixture is a combination of two or more substances which retain their own chemical identities • If you put M&Ms, Resee’s Pieces and Skittles together in a bowl, they are now a mixture, but they each retain their own identities, properties and delicious goodness

3. Mixture Types • Heterogeneous mixtures: the mixing is not uniform, and different regions of the mixture have different compositions • Homogenous mixtures: the mixing is uniform and the composition is the same throughout the mixture • Homogenous mixtures can be further categorized

4. Homogenous Mixtures • Solutions: particles are about the size of a molecule, about 0.1-2 nm • Solutions are the most important class • Colloids: particles are bigger, 2-500 nm • Milk is a colloid • Suspensions: are between fully heterogeneous an homogenous mixtures with particles bigger than 1000 nm • Many medications are suspensions • They say “shake well before use”

5. Mixture Differences • Solutions: transparent, but may have color • Colloids: can be somewhat transparent to murky to opaque, depending on the particle size • Suspensions: will be opaque, and possibly even have visible distinct particles

6. Solutions • Solids dissolved in liquids is the most common type of solution, but any phase can be dissolved in any other phase as well • Check out Table 9.2 on page 255 • The solute is dissolved in the solvent • The solvent by definition is the component which is present in the larger amount

7. Solution Process • We talked about solubility in H2O earlier • Solubility depends on the attractive forces between solute and solvent particles • Ethanol is soluble in water because both engage in hydrogen bonding, and the strengths of the hydrogen bonds are very similar

8. Like Dissolves Like • “Like dissolves like” is the rule of thumb for solubility • Substances with similar intermolecular forces will form solutions • Polar solvents will dissolve polar and ionic solutes, and non-polar solvents will dissolve non-polar solutes • “Oil and water won’t mix”

9. Some “Mixing Among the Mixers” • Some moderately polar to non-polar substances will dissolve in water to some extent • Things like chloroform or even benzene and toluene have some very slight solubility • That is the basis for many of the tests I do at my real job!

10. Solvation • Summed up beautifully by figure 9.2 on page 256 • Solvation when water is the solvent is called hydration • There is a heat change with the process (enthalpy!) • Some dissolutions are exothermic, and others are endothermic

11. Solid Hydrates • Some ionic compounds can attach water molecules to their crystalline structure • When this happens, a solid hydrate is formed Ca2SO4.½H2O • The formula for Plaster of Paris shows that one molecule of water is shared between two calcium sulfate formula units

12. Hydrates, cont • Table 9.3 on page 257 has a list of some common hydrates • Hydrates are named as if the water molecules were an atom in a covalent compound • i.e., magnesium sulfate heptahydrate: MgSO4.7H2O • Hygroscopic compounds will pull water out of the air to hydrate themselves

13. Solubility • If two substances are mutually soluble in any proportion, they are miscible • Most substances have a finite solubility limit, meaning only so much will dissolve • An equilibrium is reached, and for every solute particle dissolving, one is reforming its natural state • Putting too much sugar in Kool-Aid

14. Saturation • A solution that has dissolved the maximum amount of solute that it can is saturated • At equilibrium (no more solute will dissolve) • A substance's solubility is given in g/100mL or g/L • NaCl is 35.8g/100mL (for water) • Solubility is different in different solvents, but water is the one that’s most important

15. Temperature and Solubility • Temperature will affect solubility • For solids dissolving in water, increased temperature will often increase solubility • Sugar in coffee • But not always • Some solubility will decrease with increased temp, and others won’t be affected much at all • Figure 9.3 on page 259 shows what’s going on visually

16. Supersaturated Solutions • Solids that are more soluble at higher temperatures can form supersaturated solutions • A solution is saturated at a high temperature, and is allowed to cool slowly • The excess should form solids, but the slow cool down prevents it from occurring • The excess solute may dramatically “crash out” forming large amounts of solid rapidly

17. Solubility of Gases • For gases, the effect is opposite • Gases are higher energy, so the increased temperature makes them too energetic to stay in solution • This is why you keep pop cold • Pressure can also affect the solubility of gases in liquids (solids and liquids don’t compress, so their solubility is largely unaffected by pressure)

18. Henry’s Law • Henry’s law states that the solubility of a gas is directly proportional to the partial pressure of the gas over the liquid Solubility α Pgas • As the partial pressure goes up, solubility goes up • C/Pgas= k (ratio is constant at a given T) • Called a “Henry constant” and compares directly gas solubilities at a given T • Table 9.4 pg 261 shows Ppartial for breath

19. O, Henry continued • Can think of P as the same as heat in an endothermic reaction and apply Lechâtlier’s principle Gas + Solvent + pressure  Solution • This is how to keep fizz in the bottle

20. Units of Concentration • Concentration is how much solute is dissolved in how much solvent • We say something is “dilute” if there’s not much dissolved and “concentrated” if there is a large amount dissolved • There are several units of concentration that you will see both in the medical field and everyday life

21. Weight/Volume % • Weight/volume percent (w/v)% • The mass (weight) of the solute in grams is divided by the volume of solvent in mililiters, and then multiplied by 100% to get a percentage • (w/v)% = (g/mL) x 100%

22. Let’s Do an Example • We have 365 grams of HCl dissolved in 2 L of water. What is the (w/v)% of HCl? • Mass in g • Volume in mL • Divide • Multiply by 100% • Answer is?

23. Volume/Volume % • Volume/volume percent (v/v)% • The volume of the solute is divided by the volume of the solvent (units usually mL) and then multiplied by 100% • (v/v)% = (mL/mL) x 100%

24. Let’s Do an Example • We have 300mL of HCl dissolved in 2L of water. What is the (v/v)% of HCl? • Volume of solute in mL • Volume of solvent in mL • Divide • Multiply by 100% • Answer is?

25. Parts per Million (or Billion) • Parts per million (ppm) and parts per billion (ppb) • Parts per million is literal: one part of something out of one million is one ppm • 1 mg of 1,000,000 mg (1kg) = 1ppm • This is my world of concentrations (ppb especially)

26. PPM and PPB • When solutions in water is concerned (which is about 99% of what you will see) the units for ppm are mg/L and the units of ppb are µg/L • The book goes through the derivation for ppm, good conversion practice! • The Worked problem on page 267 is wrong, the limit for THMs is 80 ppb for the 4 THMs combined as of 2003. Shows how much the book has been updated

27. Molarity • Molarity is the number of moles of solute dissolved in one liter of solvent Symbol: M Units: mol/L • Oh, crap, I forgot all about moles! We were done with that chapter! • This is one of the most common units for actual chemistry, but you might not see it much for nursing

28. Molarity Police • Lets do an example of a molarity problem • We have 365 grams of HCl dissolved in 2 L of water. What is the molarity (M) of the solution? • Step 1? • Step 2? • Profit? • May have to convert to L from mL

29. Dilution • Dilution is the process of reducing concentration by adding more solvent • The solute amount hasn’t changed, but solvent is increased, therefore reducing the concentration • You’ve all done this • OJ from concentrate • Mixing a drink

30. Dilution Continued • The book kinda goes backwards from specific to general • Let’s do it properly starting generally C1V1 = C2V2 • Concentration 1 times volume 1 is equal to concentration 2 times volume 2 • We’ve seen these types of relationships before, no?

31. D i l u t i o n • Any units of concentration and volume can be used, as long as they are the same on both sides of the equation • If we rearrange for C2: C2 = C1(V1/V2) • V1/V2 is a dilution factor • If V2 is double V1, then the concentration is halved

32. Molarity • Specifically M1V1 = M2V2 is seen in chemistry • Lets do an example on the board • If 1 L of a 6M solution is diluted to 3 L, what is the new molarity?

33. Electrolytes • If ions are dissolved in water and the solution conducts electricity, they are called electrolytes • The charged ions carry the electricity by being attracted to the opposite charges on the attached wire • Pure water does not conduct electricity, it is the ions dissolved in it that do

34. Strengths of Electrolytes • Strong electrolytes are completely ionized when dissolved in water (ionic cmpds) • Weak electrolytes are partially ionized (typically organic acids) • Non-electrolytes do not ionize when dissolved (like sugar)

35. Equivalents and Milliequivalents • A solution often has many ions dissolved in it ( i.e., blood and other bodily fluids) • It is impossible to assign cations to anions, because they are all mixed up • Therefore we are interested in the total number and positive and negative charges of all the ions • Enter the Equivalent (Eq)

36. Equivalents • 1 Eq = molar mass (g)/|ionic charge| • If an ion has a charge of +/- 1, one equivalent is equal to it’s molar mass • If an ion has a charge of +/- 2, one equivalent is equal to half it’s molar mass • 1 milliequivalent (mEq) = 0.001 Eq • Concentrations in the body are typically low, so mEq is more common

37. Properties of Solutions • Properties of solutions are often very similar to the properties of the solvent • One difference is that solutions tend to have higher boiling points than pure solvents, and lower freezing points than pure solvents • These are called colligative properties, and depend only on the amount of particles dissolved, not the type of particle

38. Colligative Properties • Vapor pressure lowering • Boiling point elevation • Putting salt in water to boil • Freezing point depression • Salting ice on the sidewalk • Making ice cream • Antifreeze in your car is an example of both boiling point elevation and freezing point depression

39. Osmosis • When a solution and a pure solvent (or a solution with a different concentration) are separated by a semi-permeable membrane, solvent molecules pass through the membrane to the side that is less dilute (solvent is usually water) • Water is more concentrated on one side, so the rate from that side is initially faster • The sides want to come to equilibrium and have the same dilution on both sides

40. Osmotic Pressure • Because the solute can’t cross the membrane, something else causes equilibrium • As the water increases on the more concentrated side, it creates a pressure • Once the pressure reaches a level where additional water cannot come in readily, the incoming and outgoing rates equalize, and equilibrium is attained

41. Osmotic Pressure cont • The pressure at which equilibrium is attained is the osmotic pressure • Osmolarity (osmol) is the number of moles of dissolved particles per liter (mol/L) • Osmosis is a shade bit important, as all cell membranes in the body are semi-permeable membranes

42. Gin and Tonics • Blood plasma is about 0.3 osmol and is isotonicwith the cell contents • Any solution that is physiological has the same osmolarity as blood plasma • A solution that has a lower osmolarity is called hypotonic • A solution that has a higher osmolarity is called hypertonic • Check out figure 9.12 on page 281

43. Dialysis • Dialysis is like Osmosis, but the pores in the membranes are slightly bigger, so some small solute particles can pass through the membrane, but not big solute particles • This is how the kidneys work • The small nasty solutes the body wants to get rid of pass the membrane and are excreted

44. We done just did finish Ch 9 • This lecture had more slides than even that dreadful chapter 6! • Homework: I wrote my own because the book’s were a little too hard I thought. • The problems are on the following slides

45. Problem 1 • What is the weight/volume % concentration of a solution containing 300 mg of ibuprofen in 100 mL?

46. Problem 2 • What is the molarity of a solution when 117 g of NaCl is dissolved in 2 L of water?

47. Problem 3 • Which is a stronger electrolyte, KCl or sucrose, and why?

48. Problem 4 • You mix a drink by adding whiskey to pop resulting in 250 mL of drink with a concentration of 40% (v/v). Because you are a wuss it is too strong, so you dilute it to 500 mL. What is the new concentration (v/v)% ?

49. Problem 5 • Both NaCl and CaCl2 are used for salting roads. Which is more effective in lowering the freezing point of water, and why?