Define the terms heterogeneous and homogeneous .

1. Warm up • Define the terms heterogeneous and homogeneous. • What do you think is the difference between dissolving and melting?

2. Warm up • Heterogeneous: not uniform in composition • Parts are usually visually distinct • Homogeneous: uniform in composition • Parts are not visually distinct due to being well mixed at the microscopic or submicroscopic level

3. Warm up • Melting: one substance changing phase from solid to liquid • Dissolving: two substances mixing to form a solution

4. Solutions Chapter 15 (write the red!)

5. Objectives • Distinguish between homogeneous and heterogeneous mixtures and explain how dissolving is different from melting. • Define solute, solvent, solvation, dissociation, electrolyte, and aqueous. • Identify and compare the nine different solute-solvent combinations. • Compare solutions, suspensions, and colloids.

6. Defining solutions • A little review! Chemists often work with mixtures. • Mixtures consist of two or more substances, physically combined, each of which retains its properties. • Mixtures tend to be • Parts of heterogeneous mixtures are too large to mix uniformly. HETEROGENEOUS.

7. Mixtures • When particles making up a mixture are small, they can be uniformly mixed or intermingled. • There is a(n) _________ relationship between particle size and uniformity of a mixture. • When the particles are small enough and thoroughly mixed, the result is a HOMOGENEOUS mixture. inverse

8. The terminology of solutions • A solution is a homogeneous mixture – one substancedissolved in another. These two parts have special names: • Solute: the substance being dissolved. Generally, the solute is present to the lesser extent; if the two substances were originally in two different phases, the solute is the one that changes phase. • Solvent: the substance in which the other is dissolved, and is generally present to a greater extent.

9. The terminology of solutions • The process of dissolving (ordissolution) also has two “parts”: • Solvation: the process by which solvent molecules are attracted to and associate with solute molecules • Dissociation: the process by which an ionic compound splits into its component ions **Do not confuse dissociation with ionization!

10. The terminology of solutions • Example: salt dissolving in water • Solute: salt (NaCl) • Solvent: water • Solvation: water molecules being attracted to and surrounding salt ions (salt is solvated by water) • Dissociation: salt ions breaking apart from one another (the salt dissociates)

11. The terminology of solutions • Some solutions conduct electricity • Electrolytes: compounds that conduct electricity in aqueous solution OR in the molten state. • All ionic compounds are electrolytes because they dissociate into ions, which carry charge. • A compound that does not conduct electricity when either aqueous or molten is a nonelectrolyte • Sugar is a covalent compound, and does not form ions when dissolved.

12. The terminology of solutions • If a solution is made in which the solvent is water, it is known as an aqueous solution.

13. The terminology of solutions • When the solvent in a solution is an alcohol, it is known as a tincture. A tincture of iodine is a solution of iodine (solid) in alcohol.

14. The terminology of solutions • Two liquids are “miscible” when they can mix in all proportions and form a homogeneous solution (ex: water and milk) • Opposite: immiscible.

15. There are 9 types of solutions Alloys (steel, brass, etc.) Kool Aid, salt water Smoke Dental Amalgam Antifreeze, rubbing alcohol Fog Lava Coke, 7 Up, Pepsi Air

16. Careful! Solutions should not be confused with suspensions, colloids, or emulsions

17. Comparison Atoms, ions, small molecules No scattering Stable, no separation Particles not retained on filters Homogeneous Kool Aid

18. Suspensions • Finely ground particles (larger than 100 nm), when placed in a solvent, can become “suspended” • Usually visible to the naked eye (scatters light) • Will settle out due to gravity in time • Can be purified byfiltration • Considered heterogeneous

19. Colloids • A subset of suspensions • Particles or molecules too small to be seen with ordinary microscopes (between 1 and 100 nm) become suspended in the solvent; scatters light, cannot be filtered • Under the influence of gravity, colloidal particles may take months, years, or even centuries to settle out. • Also considered heterogeneous.

20. Emulsions • Colloids whose particles are in the liquid phase.

21. Comparison

22. Logistical questions • When we make and use solutions, there are some logistical questions that need to be answered: • Will the solute dissolve in the solvent? • How do you measure the concentration of a solution? How do you calculate the amount of solute and solvent to combine? • How do you change the concentration of a solution? • How do you make something dissolve faster? • How can you get more of a solute to dissolve in the same amount of solvent?

23. How do you know if something WILL dissolve? • Watch the video again • Salt dissolves in water because they are both polar. If water were not polar, or if salt were not polar (ionic), the water molecules would not be attracted to the salt ions. • This is true of all substances; polar things dissolve in polar substances, and nonpolar things dissolve in nonpolar substances. • Like dissolves like.

24. Like dissolves like • Will these dissolve? • CuCl2 in water • NaCl in CBr4 • Why are water and oil immiscible? • Water is polar, oil is nonpolar. • Why does soap dissolve in water? • Lipids form micelles, surrounding nonpolar items with polar ends pointing outwards, allowing them to be washed away. Yes! Both are polar. No! Salt is polar, CBr4 is not.

25. Homework • Work on your research project papers; have the introduction and methods sections drafted by end of the month. We will peer-review around then.

26. Logistical questions • When we make and use solutions, there are some logistical questions that need to be answered: • Will the solute dissolve in the solvent? • How do you measure the concentration of a solution? How do you calculate the amount of solute and solvent to combine? • How do you change the concentration of a solution? • How do you make something dissolve faster? • How can you get more of a solute to dissolve in the same amount of solvent?

27. How do we measure solution concentration? • What exactly IS concentration? • Concentration is measured as the amount of solute per amount of solvent or solution. • For a liquid solution, we most commonly measure molarity, but there are many other units we can use.

28. Comparing concentrations • Concentrated: A LOT of solute is dissolved in the solution. • Dilute: A little solute is dissolved in the solution. • Note: these terms are very ambiguous! They are usually used in the relative sense – one solution is more dilute or concentrated than another.

30. Molarity • Molarity is the most common way to measure the concentration of a solution. • Molarity is expressed in M, which is equal to the moles of solute per liters of total solution. • A “3 molar (or 3 M) solution” of HCl is a solution that contains three moles of HCl per liter.

31. Molality • Molality (m) is calculated as moles of solute per kilograms of solvent. • A solution containing 3 moles of solute per 1 kg of solvent is a “3 molal solution”.

32. Mole Fraction • The mole fraction of a substance is measured as the moles of solute divided by the total moles of all parts of the solution. Note, there is no unit for this quantity.

33. Mass Percent • The mass percent is measured as mass of the solute divided by the total mass of the solution, multiplied by 100. • Note that since the units will cancel out, it does not matter what unit of mass you use – as long as they are the same.

34. What if you need to change the concentration? • Many solutions (like acids) are sold and shipped at high concentration for efficiency, and need to be diluted by the buyer to lower concentrations for lab work. • Let’s say you start with 50. mL of 16 M HCl. How many moles of HCl are in this sample? 16 M HCl = x moles HCl x = 0.050 L x 16 M 0.050 L = 0.80 moles HCl

35. What if you need to change the concentration? • If you were to use the same number of moles of HCl to make a 1.0 M solution, what would be its volume? 1.0 M = 0.80 moles HCl x L solution x = 0.80 L = 800 mL (Sig figs: 8.0 x 102mL)

36. Dilution Calculations • The calculation for dilutions is very simple. Since you are not changing the number of moles of the solute, the product of the volume and molarity of the initial solution is equal to the product of the volume and molarity of the diluted solution. • Notice that the volumes can be in L or mL – as long as they match.

37. Warm up • Devise a procedure for determining the molarity and molality of an unknown sodium chloride solution. • Start with a known volume (in L) of the solution • Measure the mass of the solution • Boil the solution until only the salt is left • Mass the salt alone and convert to moles

38. Warm up • Devise a procedure for determining the molarity and molality of an unknown sodium chloride solution. • Molarity: divide the moles of salt by the volume of the solution • Molality: subtract the final mass of the salt from the initial mass of the solution to find the mass of water lost. Convert this into kg. Divide the moles of salt by the kg of water lost.

39. Do it! • How would you use your data to calculate the mass % and mole fraction? • Mass %: Divide the mass of the salt by the original mass of the solution, multiply by 100% • Mole fraction: convert mass of salt and mass of water lost into moles and divide moles of salt by total moles

40. Methods of dilution • So you would need to take your 50. mLHCl and add enough water to make 800 mL of the diluted solution. • Most people assume that this means you can measure out 750 mL of water and just add it to the 50 mLHCl. That is not always true because molecules do not always pack the same way.

41. Methods of dilution • The most common method of making a dilution requires a volumetric flask.

42. Warm up • Turn in your calculations from Tuesday’s mini-lab. I will award 2 EC pts to the group with the fewest errors and 1 EC pt to the runners-up. • List the ways in which you can make a solute dissolve in the solvent faster (hint: there are 4).

43. Logistical questions • When we make and use solutions, there are some logistical questions that need to be answered. We have answered: • Will the solute dissolve in the solvent? • How do you measure the concentration of a solution? • How do you calculate the amount of solute and solvent to combine? • How do you change the concentration of a solution?

44. Logistical questions • We still need to know: • How do you make something dissolve faster? • How can you get more of a solute to dissolve in the same amount of solvent? • How do you describe concentrations relative to the solute’s identity?

45. What affects how FAST something dissolves? • Watch the video again – what has to happen in order for the salt to dissolve? • Just like a chemical reaction, forming a solution is all about particle collisions! The more contact there is between the solute and solvent particles, the faster something will dissolve. • How do you get particles to hit each other more often?

46. Factors that affect the rate of solution formation • Heat • Note: Gases work differently • Surface area • Stirring (agitation) • Concentration

47. How about how MUCH can dissolve? • The solubility of a substance is an important property – it is defined as the extent to which a substance can be dissolved in a given solvent (how MUCH of it you can get to dissolve in a certain amount of solvent?) at a specific temp. • What affects the solubility of a substance?

48. Factors that affect solubility • Nature of solute and solvent (remember: like dissolves like) • Temperature (different for gases) • Pressure (only for gases)

49. Why do gases work differently? • When heating a solid or a liquid, what happens? • Particles move faster and farther apart; particles eventually move far enough apart that we observe a phase change – usually solid to liquid • The phase change itself requires extra energy

50. Heating Curves • If you heat ice and graph the temperature of the substance vs. time, the graph looks like this: