CHEM 121 Introduction to Fundamental Chemistry Summer Quarter 2008 SCCC Lecture 11

1. CHEM 121 Introduction to Fundamental Chemistry Summer Quarter 2008 SCCC Lecture 11 http://seattlecentral.edu/faculty/lcwest/CHE121

2. CHEM 121 Introduction to Fundamental Chemistry Summer Quarter 2008 SCCC Lecture 8 http://seattlecentral.edu/faculty/lcwest/CHE121

3. Earlier we defined matter as being either a mixture or a pure substance. Can anyone remember the two kinds of mixtures possible?

4. So far we have been talking mostly about pure substances. • e.g. ions, molecules etc • Most matter in natural systems is present as a mixture. • Solutions are a kind of homogenous mixture. • They consist of a larger component called the solvent and one or more smaller components called the solutes. Can you think of examples of solutions?

5. Examples of solutions include: Air: what is the solvent in air ? Nitrogen, N2 What is a solute in air? Oxygen, O2

6. Examples of solutions include: 18 ct gold: what is the solvent in 18 ct gold ? Gold, Au What is a solute in 18 ct gold? Copper, Cu

7. Examples of solutions include: Sea water: what is the solvent in sea water ? Water, H2O What is a solute in sea water? NaCl, salt

8. Some general properties of solutions include: • Solutions may be formed between solids, liquids or gases. • They are homogenous in composition • They do not settle under gravity • They do not scatter light (like muddy water) Solutions form when one substance dissolves in another.

9. Soluble substances are those that can dissolve in a given solvent. Insoluble or immiscible substances are those that cannot dissolve in a given solvent. Which of the following are soluble in water? NaCl, sugar, cooking oil, alcohol, gasoline, motor oil Which of the following are immiscible in cooking oil? NaCl, sugar, alcohol, gasoline, motor oil, water

10. The maximum amount of a given solute a solvent can dissolve is called the solubility. The solubility is dependent on the temperature and pressure. Solubility is often expressed in terms of grams of solute per mL of solvent but may have other units. When a solvent contains the maximum amount of a solute possible the solutions is said to be saturated.

11. By forming a solution at a high temperature then slowly cooling it we can form supersaturated solutions that contain more solute than in a saturated solution. These kinds of solutions are very unstable and tend to separate out the excess solute with the slightest disturbance. http://www.youtube.com/watch?v=uy6eKm8IRdI&NR=1 http://www.youtube.com/watch?v=aC-KOYQsIvU&feature=related

12. Solutions form when one a soluble solute is dissolved in a solvent. In biological systems aqueous (solutions where water is the solvent) are of particular importance. The solubility of most liquids and solids in water increases with temperature. The effect of pressure on the solubility of liquid or solid solutes in water is negligible.

13. The solubility of gases in water decreases with temperature. • Are cold carbonated drinks bubblier than warm carbonated drinks? The solubility of many gases in water is directly proportional to the pressure being applied to the solution. i.e. double the pressure, double the solubility • What happens when the cork is removed from a bottle of champagne? • What is the origin of decompression sickness? • Anyone heard of hyperbaric therapy?

14. How do solutions form? Why do some substances leave one phase and enter the solution and others don’t? How can we use chemistry to predict solubilities? Lets first look at the formation of a solution between an ionicsolute and a polar solvent such as H2O.

15. Ionic compounds are composed of oppositely charged ions arranged in a repeating 3-d arrangement. They are held together by attractive forces between oppositely charged ions.

16. When we place an ionic solid in water there will be attractive forces between the ions at the surface of the crystal and the water molecules. Water molecules orient such that the -ve end of the molecule is oriented towards the +ve ions at the surface and vice versa. e.g. for KCN red is the region where electrons are found most oftenand blue is where electrons are rarely found

17. H2O K+ If the attractive force between the surface ion and the solvent is greater than the forces between the ion and the solid then the ion will enter the solution phase. The ion that has left the solid and becomes completed surrounded by water molecules. It has become solvated or hydrated.

18. The process continues as new water moleculesapproach the crystal until the crystal has been fully dissolved. Note the different orientation of water molecules around the oppositely charged ions.

19. In a solution of an ionic compound a solvated ion will occasionally collide with the surface of the solid. Sometimes when this happens the ion will “stick” to the surface and become part of the solid phase again. This will happen more frequently the more concentrated the solution is. When the rate of ions leaving the solid equals the rate of ions going back to the solid the system is at equilibrium and the solution is saturated.

20. When a solution is at equilibrium with its solute macroscopically there will be no change occurring. However, at the molecular level lots is happening, just in equal and opposite directions. Supersaturated solutions can form because there are no sites for solute ions to collide with.

21. When we place a “seed” crystal in a supersaturated solution this provides the needed sites and the excess solute crystallizes very quickly. In the you tube video we watched you can just see the tiny seed crystals on the persons finger.

22. Polar but non-ionic solutes dissolve in water via a similar mechanism as for ionic compounds.

23. A solute will be insoluble in a solvent if: • Forces between solute particles are greater than the forces between solute particles and the solvent.

24. A solute will be insoluble in a solvent if: • Forces between the solvent particles are stronger than forces between the solvent and the solute. • e.g. The only attractive force between oil and water will is dispersion forces. These are weak compared to hydrogen bonds between water molecules. In a polar solvent there will be attraction between the oppositely charged ends of the molecule.

25. A good “rule of thumb” that works especially well for non-ionic compounds is: “Like dissolves like” i.e.Polarsolvents dissolve polar solutes well and non-polar solvents dissolve non-polar solutes well.

26. There are some more specific rules that allows us to better estimate the solubility of ionic compounds. You will be given these if you need them.

27. The rate of dissolution is dependent upon: • The surface area of the solute. • i.e. how finely divided it is. Increasing rate

28. The rate of dissolution is dependent upon: • How hot the solution is. • i.e. the kinetic energy of solute and solvent. • The rate of stirring. Typically when we are preparing a solution in the lab we will both heat and stir.

29. When a solute dissolves in a solvent heat can be released or absorbed. When heat is absorbed the process is endothermic and the solution becomes cooler. This effect is used in instant cold packs for sporting injuries and first aid.

30. More commonly dissolution is an exothermic process and heat is released when a solute is dissolved. Sometimes when we make a solution it will get so hot it boils!! What is the safest way to prepare a solution?

31. Many reactions take place in solution: e.g. 2AgNO3(aq) + Na2CO3(aq)  Ag2CO3(s) + 2NaNO3(aq) The co-efficients in equations allow us to determine the relative amounts of products and reactants. # moles of Ag2CO3 = ½(# moles of AgNO3) # moles of Na2CO3 = ½(# moles of AgNO3) # moles of NaNO3 = # moles of AgNO3

32. The amount of solute in a given amount of solution is defined by the concentration. The solution concentration can be defined in a variety of ways. The most useful is the molarity (M). The molarityof a solution is defined as: “The number of moles of solute in 1 L of solution” and is given the formula: M = n/V

33. If I wanted to prepare 0.5 L of 1 molL-1 NaCl aqueous solution how would I do it? M = 1 molL-1 V = 0.5 L n = ? M = n/V n = MV 1molL-1 x 0.5 L = 0.5 moles of NaCl required m = n x MW = 0.5 mol x (22.99 + 35.45) gmol-1 m = 29.22 grams of NaCl required. To make solution take 29.22 g of NaCl dilute to a volume of 0.5L with H2O

34. In the lab we would use a piece of glassware called a volumetric flask to prepare this solution.

35. An alternative way of expressing the concentration of a solution is in percent. Which is defined as: “parts of solute per 100 parts of solution” or: • Three different percent concentrations are commonly used: • % weight/weight (w/w) • % volume/volume (v/v) • % weight/volume (w/v)

36. %(w/w) is calculated using the formula: %(w/v) is calculated using the formula: %(v/v) is calculated using the formula:

37. Often we will want to make a dilute solution from a more concentrated one. To determine how to do this we use the formula : C1V1 = C2V2 Where: C1 = concentration of more concentrated solution V1 = volume required of more concentrated solution C2 = concentration of more dilute solution V2 = volume of more dilute solution We can use any units in this equation but they must be the same on both sides.

38. How would I prepare 50 mL of a 10 mgL-1 solution of NaOH using a 100 mgL-1 stock solution? C1V1 = C2V2 V1 = C2V2/C1 V1 = (10 mgL-1 x 50 mL)/100 mgL-1 V1 = 5 mL Take 5 mL of stock NaOH solution and dilute to 50 mL