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Chapter 13: Solutions

Chapter 13: Solutions. The Classification of Matter. Solutions. Solution : a homogenous mixture of two or more substances in a single phase; a solute and a solvent . Solute : a substance which is the portion of a solution that is dissolved.

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Chapter 13: Solutions

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  1. Chapter 13: Solutions

  2. The Classification of Matter

  3. Solutions • Solution: a homogenous mixture of two or more substances in a single phase; a solute and a solvent. • Solute: a substance which is the portion of a solution that is dissolved. • Solvent: a substance which is the portion of a solution that dissolves the solute; determines the phase of the solution.

  4. How does a solvent dissolve the solute, you might ask? • The solvent has enough energy to break the surface tension or the bonds of the solute. • Remember those pesky intermolecular forces?! • The strongest of all are the metallic bonds followed by the hydrogen, ion-dipole, dipole-dipole, and finally the London dispersion forces.

  5. Phases of Solutions

  6. Suspensions • Suspension: A mixture whose particles are evenly dispersed in a gas or a liquid and which settle out over time. Examples: sand – water mixture, cement

  7. Colloids • Colloid: amixture whose particles are smaller than a suspension but larger than a solution and does not settle; the particles are suspended in a liquid, gas or solid. • The Tyndall effect is a phenomenon where a beam of light is reflected off of the particles and visible to the naked eye. • Examples: milk, mayonnaise, marshmallow, fog

  8. Separation Anxiety • One difference in mixtures and compounds is that mixtures can be separated into it’s parts. Mixtures can be separated by chromatography, gel electrophoresis, distillation, mechanical separation, evaporation, etc. Can you think of any more?

  9. What happens to a suspension when it is allowed to stand over a period of time? • It’s components separate into separate layers

  10. Concentrate on This • Concentration is the measure of the amount of a particular substance in a given volume of solution.

  11. Concentration Calculations

  12. Molarity • Molarity: The molar unit is probably the most commonly used chemical unit of measurement. Molarity is the number of moles of a solute dissolved in a liter of solution. A molar solution of sodium chloride is made by placing 1 mole of a solute into a 1-liter volumetric flask. Water is then added to the volumetric flask up to the one liter line.

  13. 1M NaCl in water • The result is a one molar solution of sodium chloride. • M = n/V • (remember that n = mass / molar mass) • M = 1 mol NaCl / 1 L H2O & • 1 mol NaCl = 58.44 grams NaCl • So,M = 58.44 grams NaCl / 1 L H2O

  14. Molality • Molality: The molal unit is not used nearly as frequently as the molar unit. A molality is the number of moles of solute dissolved in one kilogram of solvent. Be careful not to confuse molality and molarity. Molality is represented by a small "m," whereas molarity is represented by an upper case "M."

  15. Notes on Molality • Note that the solvent must be weighed unless it is water. One liter of water has a specific gravity of 1.0 and weighs one kilogram; so one can measure out one liter of water and add the solute to it. • Most other solvents have a specific gravity greater than or less than one. • Therefore, one liter of anything other than water is not likely to occupy a liter of space.

  16. 1m NaCl(aq) • To make a one molal aqueous (water) solution of sodium chloride (NaCl) , measure out one kilogram of water and add one mole of the solute, NaCl to it. • The formula weight for NaCl is 58, and so 58.44 grams of NaCl dissolved in 1kg water would result in a 1 molal solution of NaCl. • m = 1 mol NaCl / 1 kg H2O • m = 58.44 g NaCl / 1 kg H2O

  17. Normality • Normality: There is a relationship between normality and molarity. Normality can only be calculated when we deal with reactions, because normality is a function of equivalents. Normality is the molarity of a solution multiplied by the number of moles of equivalents. • Normality is particularly useful in titrations calculations.

  18. % by Weight • Percent by weight: To make up a solution based on percentage by weight, one would simply determine what percentage was desired (for example, a 20% by weight aqueous solution of sodium chloride) and the total quantity to be prepared.

  19. If the total quantity needed is 1 kg, then it would simply be a matter of calculating 20% of 1 kg which, of course is: 0.20 NaCl * 1000 g/kg = 200 g NaCl/kg. • In order to bring the total quantity to 1 kg, it would be necessary to add 800g water.

  20. % by Volume • Percent by volume: Solutions based on percent by volume are calculated the same as for percent by weight, except that calculations are based on volume. Thus one would simply determine what percentage was desired (for example, a 20% by volume aqueous solution of sodium chloride) and the total quantity to be prepared.

  21. If the total quantity needed is 1 liter, then it would simply be a matter of calculating 20% of 1 liter which, of course is: 0.20 NaCl * 1000 ml/l = 200 ml NaCl/l. • Percentages are used more in the technological fields of chemistry (such as environmental technologies) than they are in pure chemistry.

  22. Making Dilutions • Making dilutions is easy. The starting solution is referred to as the stock solution and the solution that you make is referred to as theworking solution. • It is easier to make and store smaller volumes of concentrated stock solution.

  23. Calculating Dilutions • When you need a reagent (A substance used in a chemical reaction to detect, measure, examine, or produce other substances;) simply prepare it from a stock solution. • The molarity and volume of the stock solution are inversely proportional. As you reduce one, you increase the other by the same proportion. Therefore; • MstockVstock = MworkingVworking

  24. You Calculate! • For example: You start out with a 12M concentrated solution of saline solution. The experiment calls for 0.8 Liters of a 3.0M solution. How would you prepare it? • 12M x Vol. Stock = 3.0M x 0.8L • Vol. Stock = 3.0M x 0.8L / 12M = 0.2L • Therefore, you take 0.2L of your stock solution and dilute to 0.8L final volume.

  25. Soluble or Insoluble • What makes a substance soluble? Solubility is the relative ability of a substance to dissolve into another at a given temperature and pressure; or the amount of solute that will dissolve into a given solution to saturate the solution.

  26. Miscible or Immiscible • Miscible: The tendency or capacity of two or more liquids to form a uniform mixture, that is, to dissolve in each other • Immiscible: Liquids are that do not mix (e.g. oil and water)

  27. Measuring Solubility • Solubility is measured in g solute / 100 mL solvent • For example, if a solid powder has a solubility of 10g/100 mL of water, then 1.0 grams of the same solid would need 10 mL of water to dissolve.

  28. Saturated or Unsaturated • A saturated solution cannot dissolve any more solute – like a glass of iced tea that you added too much sugar to and the sugar settles on the bottom. An unsaturated solution contains less solute that the saturated solution – like the perfect glass of sweetened iced tea, with no settled solute (sugar) on the bottom.

  29. Supersaturated • A supersaturated solution holds more dissolved solute than what is required to reach solubility equilibrium at a given temperature - or a solution that contains more solute than it would if the dissolved solute were in equilibrium with the undissolved solute.

  30. Equilibrium • A solubility equilibrium is the state at which a solution is in equilibrium with its solute ions and the solvent. • SO, try this at home. Saturate a pan of water with sugar, add more sugar so that it settles on the bottle. Heat the water. What happens to the excess sugar? (It dissolves into the water.) Allow the water to cool. Did the excess sugar fall out of solution? (No) Now, that is a supersaturated solution!

  31. Effects on Solubility • The speed of solubility depends on the intermolecular forces, mass, surface area, temperature. As temperature increases, the solubility of most substances increases

  32. Solubility Curves • The effects of an increase in temperature on selected compounds is shown. • You can determine the solubility of select compounds at different temperatures.

  33. Interpreting the Curves • To the right you see the solubility curve for KNO3 • As you cool a 100 g KNO3 per 100 g H2O solution,the solid will appear first at 57oC

  34. Solute on Solvent • What effects do a solute have on a solvent? To answer that question, consider what happens when you apply salt (NaCl) to an icy road. Does the salt melt the ice? • No, what the salt does is lower the melting or freezing point of water .

  35. What about the Boiling Point? • What happens when you add salt to a pot of water on the stove? • It raises the boiling point of the water.

  36. Colligative Properties • So using this example, a solute (the salt) decreases the melting/freezing point of the solvent (water) and raises the boiling point. • The melting/freezing and boiling points of a solution are colligative properties. • ex. Vapor pressure, boiling point, freezing point, osmotic pressure

  37. What does that mean? • A colligative property is a property of a solution that is dependent upon the number of solute particles present and the nature of the solvent. • These properties are independent of the identity of the solute. • In other words, if you replaced 1 mole salt with 1 mole of sugar in the previous example, the effects on mp/fp and bp would be the same!

  38. The greater the concentration of solute, the greater the effect it has on a colligative property. • A solute also lowers the vapor pressure of a given solvent.

  39. Calculating BP and MP • The change in the bp and mp can be calculated for solutions if the solute is molecular (covalent bonding, not ionic) and non-volatile (does not evaporate) because the bp and fp are proportional to the molality.

  40. Calculating Temperature • Tb = (+kb) x (m) and; Tf = (-kf) x (m) where “+” indicates and increase and “-“ indicates a decrease in temperature. • The symbols kb and kf are the boiling point and freezing point constants (oC/mol) and represents the number of degrees in Centigrade that the bp or fp is raised or lowered when 1 mole of a molecular, non-volatile solute is dissolved in 1 kg of a solvent.

  41. Henry’s Law • Henry’s Law states that at a constant temperature, the solubility of a gas is directly proportional to the partial pressure of the gas on the surface of the liquid. • What happens when you open a soda?

  42. Open it Up • When you open a soda, either in a bottle or in a can, there is a sudden release of pressure and the gas dissolved on the surface of the soda, comes out of solution (it foams up.) • Soda is a saturated solution of CO2 and water, sugar, etc. The carbon dioxide is exerting partial pressure on the surface of the soda.

  43. Carbonated Beverages • Soda has a high pressure when it is first bottled or canned. This high pressure environment increases the solubility of the CO2; therefore, more carbon dioxide is dissolved into the soda. • When the pressure is released, when the soda is opened, the pressure in the container equals the pressure of the atmosphere, forcing the excess CO2 to come out of solution.

  44. Electrolytes • Sometimes, solutions exhibit different characteristics than the parent solute or solvent. • (Does that sound familiar? How about mother and daughter ions!) • An electrolyte is a solute that is able to conduct electricity when dissolved in a solvent.

  45. Conductivity • It is said to have conductivity. Conversely, a non-electrolyte does not conduct electricity in solution or out of solution. Ions can be produced in solution 2 ways: • .An ionic compound is separated into it’s ions by a process called dissociation; or • .A polar covalent compound dissolves in water and loses a Hydrogen to form a Hydronium ion: HCl(g) + H2O(l) H3O+(aq) + Cl-(aq)

  46. Effects • The effect of these electrolytes on the bp and fp are higher or lower than expected based on the k x m equation. Why? • Because, 1 mole of an ionic compound produces more than 1 mole of electrolytes! Let’s see this in action. One mole of NaCl produces 1 mole of Na+ and 1 mole of Cl- for a total of 2 moles of electrolytes. Therefore, the effect on the bp/fp is 2 times the effects of a molecular substance.

  47. Clean it Up! • What do you wash your clothes with? In Chemistry terms, laundry detergent is referred to as a surfactant. • A surfactantor surface active agent, is a molecule that is amphiphilic which means that it has a water loving end (hydrophilic) and water fearing end (hydrophobic).

  48. Surfactants • These agents are wetting agents that lower the surface tension of a liquid, allowing easier spreading, and lower the interfacial tension between two liquids.

  49. Detergents and Soap • Detergent is a manufactured water soluble surfactant that emulsifies oil and dirt; makes a suspension from one liquid into another where the first liquid does not dissolve in the second (kind of like oil and vinegar salad dressing). • Soap is a naturally occurring detergent that emulsifies dirt.

  50. Classwork and Homework • Chapter Review, page 488 • 2, 7-8, 17-18, 25-28, 44-45, 51-52,61-63

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