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Solutions

Solutions. Solution Basics. Solution = homogeneous mixture in which a solute is dissolved in a solvent Not limited to solids dissolved in water Solutions can be solids, liquids or gases. Key Terms. Solute- the material being dissolved Solvent-the dissolving medium

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Solutions

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  1. Solutions

  2. Solution Basics • Solution = homogeneous mixture in which a solute is dissolved in a solvent • Not limited to solids dissolved in water • Solutions can be solids, liquids or gases

  3. Key Terms • Solute- the material being dissolved • Solvent-the dissolving medium • Solution- homogeneous mixture • Soluble- capable of dissolving (sugar is soluble in water) • Insoluble- not capable of dissolving (lead sulfide is insoluble in water • Miscible- fluids that mix (liquid,gas) • Immiscible- fluids that don’t mix (oil and water are immiscible in each other) • Note: specify solute and solvent: It is not enough to say salt dissolves. It does dissolve in water, but not in hexane.

  4. The dissolving process • The process by which an ionic compound dissolved in water involves water molecules surrounded (solvating) the ions. Because the water molecules have charges (polar molecule) they can interact with the ions.

  5. Why is the O end of the water molecules “facing” in toward the Na+ ion? But the H end faces In toward the Cl- ion? Na+ Cl-

  6. What determines how well something dissolves? • There must be an interaction between the molecules of solute and solvent • There is also a tendency for substances to mix regardless of interactions (this is related to a concept known as entropy which we’ll discuss later)

  7. Ionic Materials • Tend to dissolve in polar solutes (water) as the ions can interact with the solute molecules • The sum of the interactions must overcome the electrical attraction between ions (lattice energy). • Lattice energy depends on charge and size.

  8. Lattice Energy • Electrical attraction depends on charge and distance (F = q1q2/d2) • Ions with +1 or -1 are easier to separate than those with charges of 2 or 3 (+ or -). That is why the soluble ions are all +1 or -1 (solubility rules: Na+1, K+1, NO3-1, NH4+1 and C2H3O2-1 are almost always soluble)

  9. Size • Here it gets complicated • Larger ions are separated by a greater distance so their attractions are less, making them tend to be more soluble • But, smaller ions have more charge concentration (nature does not like a lot of charge or energy concentrated in one place) which would make small ions more soluble • Which factor matters more (it varies a lot and it depends) That’s a big help, I know.

  10. Molecular Compound Solubility • In order for two liquids to mix with each other there be some interactions between the molecules. • Two polar liquids can mix with each other (be miscible). • Two nonpolar liquids can be miscible in each other. • A polar liquid and a nonpolar liquid will not be miscible

  11. Why • Essentially, when molecules mix with each other they are “trading off” interactions • Water can form H-bonds with water • Methanol can form H-bonds with methanol • Water can form H-bonds with methanol • It is okay for water to interact with methanol since the H-bond interaction is equally “favorable” whether it is between two waters or a water and a methanol • Plus, there is an added “bonus” from mixing and increasing entropy

  12. Like Dissolves Like • In general, polar substances dissolve well in other polar substances (think of ionic as extremely polar) • Nonpolar substances dissolve well in nonpolar substances

  13. Kitchen Examples Water salt sugar vegetable oil (polar) (ionic) (polar) (nonpolar) What dissolves in water (salt and sugar) What does not (oil) Will salt or sugar dissolve in oil? Why? Ammonia dissolves in water so it must be… Iodine dissolves in oil, so it must be…

  14. Oil and Water • They don’t mix because the water molecules are all so “busy” attracting each other with hydrogen bonds • An interaction between polar water and nonpolar oil could involve London forces which are less energetically favorable. • In a sense, the water won’t trade off hydrogen bonds for London Forces • They don’t “repel” each other

  15. Examples of Solubility

  16. A Pattern For Illustration • The greater the hydrogen bonding capability of a molecule, the more water soluble (or miscible) it will be. • Alcohols (with an –OH) are less soluble as the molecule gets larger (the OH is a smaller percentage of the molecule) • Molecules with two OH groups (glycols) are more soluble

  17. Solubility Curves- Solids • Graphs show the limits of how much solute can be dissolved per 100g solvent at a given temperature • The line shows what would make a saturated solution • Below the line (more solute can be dissolved) = unsaturated solution • Above the line (more solute dissolved than is supposed to happen) = supersaturated solution

  18. Most solids (not all) are more soluble as T increases)

  19. How many grams of KBr can dissolve in 100 g water at 60°? 85 grams How many grams of KBr can dissolve in 200 g water at 60°? Twice as much solvent can dissolve twice as much solute = 170 grams If 60 grams of NaClO3 are placed in 100 grams of water at 40°C, what type of solution is made? Unsaturated (below the line which sows a saturated solution)

  20. Solubility Curves- Gases • Gases are more soluble as temperature decreases and as pressure increases • Solubility of solids is not really affected by P (only T) • Easy to remember- soda stays “fizzy” in the fridge (low T) and with the cap on (higher P)

  21. Temperature Effects • Two points: • As T increases, solubility • Decreases • b) Solubility is much lower • for gases than solids (look • at y-axis) How does T affect aquatic life? (thermal pollution) Why do active fish like colder water?

  22. Pressure Effects • Gas solubility is directly related (proportional) to Pressure • This is known as Henry’s law

  23. Solution Concentration Descriptions • Here we (we means you like when your parents used to say “we’re going to the dentist”) learn ways to describe solution concentration and how to do conversions between these units • It helps to keep in mind exact definitions of solute, solvent and solution • Remember that density = mass/volume and that mass = volume X density

  24. Concentration Units • Mole Fraction • Mass Percent • Molarity • molality

  25. Mole Fraction (X) • Ratio of moles of substance x over total moles in mixture • Has no units (moles/moles = dimensionless) • Example: What is the mole fraction of each component in a solution of .45 moles of benzene and .65 moles of toluene • (it does not matter if you know what benzene and toluene are)

  26. Mole Fraction #1 • Xben = .45/1.1 = .41 • Xtol = .65/1.1 = .59 • Sum of mole fractions = 1

  27. Example #2 • Example #2: What is the mole fraction of sucrose (C12H22O11) in a solution made by dissolving 171 g sucrose in 180 grams of water. • Xsucrose= moles sucrose/total moles 0.5 moles sucrose/(10.0 moles water +0.5 moles sucrose)= .5/10.5 = .048

  28. Mass % • Mass % solute = grams solute/mass of solution • Keep in mind that solution = solute + solvent • An isotonic solution of NaCl is 0.9%, which means that for each 100g solution, there are 0.9 g NaCl

  29. Mass % Example 1 • How many grams of KCl are needed to make 525.0 g of 3.10% KCl • Out of the 525g solution, there must be: (0.0310)(525.0) g KCl = 16.3 g KC • The rest (525.0-16.3 = 508.7g) is the mass of water • You add 16.3 g KCl to 508.7 g water.

  30. Mass % Example 2 • In an experiment you need 42.0 g of an HCl solution that is 18.0% HCl. • How many grams of HCl are used? • Mass of HCl = (42g)(.18) = 7.6 g HCl

  31. Molarity (M) • Moles of solute/L solution • Most commonly used unit of solution concentration in chemistry (and the one we have already done). • Example 1: What is the molarity of a solution made by dissolving 35.6 g KBr in a solution volume of 400 ml?

  32. Example 1 Moles KBr = 35.6 g KBr 1 mole KBr = 119 g KBr .299 moles/.400L = .75M (molar)

  33. Example 2 • What mass of sucrose (C12H22O11) is needed to make 350 ml of .25M solution. M= moles solute L solution .25M = x moles .350L X = .0875 mole x 342 g/mole = 29.9 g sucrose

  34. Volumetric Flask Glassware used for accurately making solutions (see molarity notes from ch 4 for more detail) 1 mole in 1 L = 1 Molar solution

  35. Molality (m) • Moles of solute per kilogram of solvent • Note two differences with molarity- mass instead of volume and solvent instead of solution. • Used because mass is not affected by Temperature, while volume is

  36. Molality example 1 • What is the molality of a solution made by dissolving 67.8 g HF in 780 g water? • m = moles HF Kg water (the solvent) = 67.8 g HF 1 mole HF = 20.0 g HF .780 Kg 4.35m (molal)

  37. Molality example 2 • What mass of CaCl2 must be dissolved in 500 g water top make a .5m solution • m = moles calcium chloride/Kg water • Moles calcium chloride = .5m x .5 Kg = • .25 moles CaCl2 110 g = 27.5 g CaCl2 1 mole CaCl2

  38. Conversions Between Concentration Units • Keys- the one thing that ties volume and mass together is density • Solution = solute + solvent • Label everything (unit and id as solute, solvent or solution as appropriate • We’ll look at several examples and you’ll get lots of practice

  39. molality  mole fraction • An aqueous solution of sucrose (C12H22O11) has a molality of 0.150m. What are the mole fractions of sucrose and water in the solution? • 0.150 m means 0.150 moles of sucrose for each 1 kg of water. That ratio holds no matter what the exact amounts are. 1 kg of water = 1000 g = 1000 g water 1 mole water = 18.0 g water 55.6 moles water.

  40. molality  mole fraction cont’d • Mole fraction of sucrose = moles sucrose = total moles = 0.150 moles/(0.150 moles + 55.6 moles ) = .00269 • Mole fraction of water = moles water = total moles = 55.6 moles/(0.150 moles + 55.6 moles) = 0.997

  41. Mole fraction  molality • A solution of sodium bromide has a NaBr mole fraction of 0.250 and a water mole fraction of 0.750. What is the molality of the NaBr in the solution? • The molality is = moles NaBr divided by Kg water. We already know there are .250 moles of NaBr and there would be • 0.750 moles H2O 18g H2O 1 Kg = 1mole H2O 1000 g

  42. Mole fraction  molality cont’d • Molality = .250 moles NaBr = 0.0135 kg H2O 18.5 m

  43. Mass %  mole fraction • What are the mole fractions in an aqueous 5.4% NaF solution? • Assume 1000 g • 5.4% of the mass (54 grams is NaF) and • 94.6% of the mass (946 g is H2O) • Mole NaF = 54 g NaF 1 mole NaF = 1.29 42 g NaF • Mole H2O = 946 g H2O 1mole H2O = 52.6 18 g H2O

  44. Mass %  mole fraction cont’d • Mole fraction NaF = 1.29/(1.29 +52.6) = .024 • Mole fraction H2O = 52.5/(1.29 + 52.6) = • .976

  45. Molality  mass % • What is the mass % of a .650 m solution of benzene (C6H6) dissolved in toluene (C7H8)? • This means that we can think of the solution as being made of .650 moles of benzene dissolved in 1000 g toluene • We could use any similar ratio • .650 moles C6H6 78 g C6H6 = 50.7 g C6H6 1 mole C6H6

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