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Chapter 16 Properties of Solutions

Chapter 16 Properties of Solutions. Solution Formation. Solutions are homogeneous mixtures that may be solid, liquid, or gaseous. The compositions of the solvent and the solute determine whether a substance will dissolve.

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Chapter 16 Properties of Solutions

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  1. Chapter 16Properties of Solutions

  2. Solution Formation Solutions are homogeneous mixtures that may be solid, liquid, or gaseous. The compositions of the solvent and the solute determine whether a substance will dissolve. Stirring (agitation), temperature, and the surface area of the dissolving particles determine how fast the substance will dissolve. These three factors involve the contact of the solute with the solvent.

  3. Stirring & Solution Formation Stirring speeds up the process of dissolving because fresh solvent is continually brought into contact with the surface of the solute Stirring affects only the rate at which a solid solute dissolves. It does not influence the amount of solute that will dissolve. An insoluble substance remains undissolved regardless of how vigorously or for how long the solvent/solute system is agitated.

  4. Temperature & Solution Formation At higher temperatures, the kinetic energy of the solvent molecules is greater than at lower temperatures so they move faster. The more rapid motion of the solvent molecules leads to an increase in the frequency and the force of the collisions between the solvent molecules and the surfaces of the solute molecules.

  5. Particle Size & Solution Formation A spoonful of granulated sugar dissolves more quickly than a sugar cube because the smaller particles in granulated sugar expose a much greater surface area to the colliding solvent molecules. The more surface of the solute that is exposed, the faster the rate of dissolving.

  6. Solubility If you add 36.0 g of NaCl to 100 g H2O at 25ºC, all of the 36.0 g of salt dissolves. If you add one more gram of salt and stir (no matter how vigorously or for how long) only 0.2 g of the last portion will dissolve. According to the kinetic theory, water molecules are in continuous motion. They should continue to bombard the excess solid, removing and solvating the ions.

  7. Solubility As ions are solvated, they dissolve in the water. Based on this information, you might expect all of the salt to dissolve eventually. That does not happen, however, because an exchange process is occurring. New particles from the solid are solvated and enter into solution. At the same time an equal number of already dissolved particles crystallize. These particles come out of solution and are deposited as a solid. The mass of undissolved crystals remains constant.

  8. Solubility In a saturated solution, a state of dynamic equilibrium exists between the solution and the excess solute. The rate of solvation (dissolving) equals the rate of crystallization, so the total amount of dissolved solute remains constant. The system will remain the same as long as the temperature remains constant. Saturated solution – contains the maximum amount of solute for a given quantity of solvent at a constant temperature and pressure.

  9. Solubility Example: 36.2 g of salt dissolved in 100 g of water is a saturated solution at 25ºC. If additional solute is added to this solution, it will not dissolve. Solubility of a substance is the amount of solute that dissolves in a given quantity of a solvent at a specified temperature and pressure to produce a saturated solution. Solubility is often expressed in grams of solute per 100 g solvent. (gas sometimes g/L)

  10. Solubility Unsaturated solution – a solution that contains less solute than a saturated solution at a given temperature and pressure. If additional solute is added to an unsaturated solution, it will dissolve until the solution is saturated. Some liquids are infinitely soluble in each other. Any amount will dissolve in a given volume. Two liquids are miscible if they dissolve in each other in all proportions (water and ethanol)

  11. Factors Affecting Solubility Temperature affects the solubility of a solid, liquid and gaseous solutes in a solvent. Both temperature and pressure affect the solubility of gaseous solutes. The solubility of most solid substances increases as the temperature of the solvent increases. Mineral deposits form around the edges of hot springs because the hot water is saturated with minerals. As the water cools, some of the minerals crystallize because they are less soluble at the lower temperature.

  12. Factors Affecting Solubility For a few substances, solubility decreases with temperature. Supersaturated solution – contains more solute than it can theoretically hold at a given temperature. Make a saturated solution of sodium acetate at 30·C and let the solution stand undisturbed as it cools to 25ºC. You would expect that solid sodium acetate will crystallize from the solution as the temperature drops. But no crystals form.

  13. Supersaturated Solutions The crystallization of a supersaturated solution can be initiated if a very small crystal, called a seed crystal, of the solute is added. Rock candy is another example of crystallization in a supersaturated solution. A solution is supersaturated with sugar and seed crystals cause the sugar to crystallize out of solution. A supersaturated solution crystallized rapidly when disturbed.

  14. Temperature and Gas Solubility The solubilities of most gases are greater in cold water than in hot. Thermal pollution happens when an industrial plant takes water from a lake for cooling and then dumps the heated water back into the lake. The temperature of the lake increases which lowers the concentration of dissolved oxygen in the lake water affecting aquatic animal and plant life.

  15. Pressure and Solubility Changes in pressure have little affect on the solubility of solids and liquids, but pressure strongly influences the solubility of gases. Carbonated beverages contain large amounts of carbon dioxide dissolved in water. Dissolved CO2 makes the drink fizz. The drinks are bottle under higher pressure of CO2 gas, which forces large amounts of the gas into solution. When opened, the partial pressure of CO2 above the liquid decreases.

  16. Pressure and Solubility Immediately, bubbles of CO2 form in the liquid and escape from the bottle and the concentration of dissolved CO2 decrease. If the drink is left open, it becomes “flat” as it loses its CO2. Henry’s Law – sated that at a given temperature, the solubility (S) of a gas in a liquid is directly proportional to the pressure (P) of the gas above the liquid. As the pressure of the gas above the liquid increases, the solubility of the gas increases.

  17. Pressure and Solubility Henry’s Law S1 = S2 P1 P2

  18. Question The solubility of a gas in water is 0.16 g/L at 104 kPa. What is the solubility when the pressure of the gas ins increased to 288 kPa. Assume the temperature remains constant. S1 = S2 P1 P2 (288 kPa) ( 0.16g/L) = 4.4 x 10-1 g/L (104 kPa)

  19. Questions What determines whether a substance will dissolve? Chemical composition of the solute and solvent. What determines how fast a substance will dissolve? Agitation, temperature and particle size of the solute What units are usually used to express the solubility of a solute? g of solute per 100 g solvent

  20. Questions What would you do to change a saturated solid/liquid solution to an unsaturated solution? Add solvent What would you do to change a saturated gas/liquid solution to an unsaturated solution? Increase the pressure What are two conditions that determine the mass of solute that will dissolve in a given mass of solvent? Temperature and pressure (if the solute is a gas)

  21. End of Section 16.1

  22. Concentration Concentration of a solution is a measure of the amount of solute that is dissolved in a given quantity of solvent. Dilute solution is one that contains a small amount of solute. Concentrated solution – contains a large amount of solute. In chemistry the most important unit of concentration is molarity.

  23. Molarity Molarity (M) is the number of moles of solute dissolved in one liter of solution Molarity (M) = moles of solute / liters of solution. Note that the volume involved is the total volume of the resulting solution, not the volume of the solvent alone. 3 M NaCl is read as “three molar sodium chloride”

  24. Molarity Questions A solution has a volume of 2.0 L and contains 36.0 g of glucose (C6H12O6). If the molar mass of glucose is 180 g/mol, what is the molarity of the solution? M = moles of solute / L of solution M = 36.0 g glucose 1 mol glucose 180 g glucose 2.0 L M = 0.1mol/L or 0.1M C6H12O6

  25. Molarity Questions A solution has a volume of 250 mL and contains 0.70 mol NaCl. What is its molarity? M = moles of solute / L of solution M = 0.70 mol NaCl 0.250L solution (convert to L) M = 2.8 mol/L or 2.8M NaCl

  26. Molarity Questions How many moles are in each quantity of the following substances? 12.0 g NaCl (12.0 g NaCl) (1 mole NaCl / 58.5 g NaCl) = 0.205 mole NaCL 53.8 g KNO3 (53.8 g KNO3) ( 1 mol KNO3 / 101.1g KNO3) = 0.532 mol KNO3

  27. Molarity Questions Find the mass in grams of each of these amounts of substances? 1.5 mol NaOH (1.5 mol NaOH) (40 g NaOH / 1 mol NaOH) = 60 g of NaOH 0.575 mol NaHCO3 (0.575 mol NaHCO3)( 84g NaHCO3 / 1 mol NaHCO3) 48.3 g NaHCO3

  28. Molarity Questions Which contains more molecules 1.00 mol SO2 or 1.00 mol SO3? They both contain the same number of molecules Which contains more mass 1.00 mol SO2 or 1.00 mol 1 mole of SO3 (80.1g) – 1 mole of SO2 (64.1g)

  29. Molarity Sometimes you may need to determine the number of moles of solute dissolved in a given volume of solution. How many moles are in 2.00 L of 2.5M lithium chloride (LiCl)? Moles of solute = molarity (M) x liters of solution (V) Moles of solute = (2.5 moles/L) ( 2.00L) Moles of solute = 5.0 mol

  30. Molarity Questions How many moles of ammonium nitrate are in 335 mL of 0.425 M NH4NO3? mol NH4NO3 = M x L of solution = (0.425 mol/L) (0.335L) = 0.142 mol NH4NO3 How many moles of solute are in 250 mL of 2.0M CaCl2? How many grams of CaCl2 is that? mol CaCl2 = (2.0 mole/L) (0.250L) = 0.50 mol 0.50 mol CaCl2 (111.11 g CaCl2 / 1 mol CaCl2 ) = 56 g

  31. Making Dilutions Diluting - To make less concentrated by adding solvent. Diluting a solution reduces the number of moles of solute per unit volume, but the total number of moles of solute in solution does not change. Moles of solute before dilution = moles of solute after dilution moles of solute = M x L of solution and total number of moles of solute remains unchanged upon dilution. M1V1 = M2V2

  32. Making Dilutions M1V1 = M2V2 molarity & volume molarity and volume of original solution of diluted solution Volumes can be L or mL as long as the same units are used for both V1 and V2

  33. Making Dilutions A student is preparing a 100 mL of 0.40M MgSO4 from a stock solution of 2.0 M MgSO4. How would she do this? M1V1 /M2 = V2 V2 = 20 ml – She would measure 20 mL of the stock solution (2.0 M MgSO4) and transfer it to a volumetric flask. Then she would add water to the flask to make 100 mL of solution. Try the class activity on page 482

  34. Questions How many milliliters of a solution of 4.0 M KI are needed to prepare a 250.0 mL of 0.760 M KI? V1 = (0.760M)(250.0 mL) / (4.0 M) = 47.5 mL How could you prepare 250 mL of 0.20M NaCl using on a solution of 1.0M NaCl and water? V1 = (0.20M) ( 250 mL) / ( 1.0 M) = 50 mL Use a pipet to transfer 50 mL of the 1.0M solution to a 250 mL flask. Then add distilled water up to the mark.

  35. Percent Solutions (v / v) The concentration of a solution in percent can be expressed in two ways: As the ratio of the volume of the solute to the volume of the solution or as the ratio of the mass of the solute to the mass of the solution Percent by volume (% (v/v)) = volume of solute x 100% volume of solution How many milliliters of isopropyl alcohol are in 100 mL of 91% alcohol?

  36. Question If 10mL of acetone (C3H6O) is diluted with water to a total solution volume of 200mL, what is the percent by volume of acetone in the solution? Percent by volume (% (v/v)) = volume of solute x 100% volume of solution % by volume of acetone = 10 mL / 200 mL =5.0% v/v

  37. Question A bottle of the antiseptic hydrogen peroxide is labeled 3.0% (v/v). How many mL hydrogen peroxide are in a 400.0 mL bottle of this solution? Percent by volume (% (v/v)) = volume of solute x 100% volume of solution 0.03 = x mL / 400.0 mL (0.03) (400.0 mL) = x 12 mL = x

  38. Percent Solutions (mass/mass) Another way to express the concentration of a solution is as a percent (mass/mass), which is the number of grams of solute in 100 g of solution. A solution containing 7 g of NaCl in 100 g of solution is 7% (m/m) Percent by mass (% (m/m) = mass of solute x 100% mass of solution

  39. Percent Solutions (mass/mass) You want to make 2000g of a solution of glucose in water that has a 2.8% (m/m) concentration of glucose. How much glucose should you use? Percent by mass (% (m/m) = mass of solute x 100% mass of solution 2000 g solution=(2.8g glucose/100 g solution) = 56 g glucose How much solvent should be used? The mass of the solvent equals the mass of the solution minus the mass of the solute. (2000 g – 56 g ) = 1944 g of solvent Thus a 2.8% (m/m) glucose solution contains 56 g of glucose dissolved in 1944 g of water.

  40. Questions How do you calculate the molarity of a solution? Molarity (M) = moles of solute / volume of solution Compare the number of moles of solute before dilution with the number of moles of solute after dilution. The total number of moles of solute in solution does not change (only the number of moles of solute per unit volume) What are two ways of expressing the concentration of a solution as a percent? Volume solute Mass of solute Volume of solution Mass of solution

  41. End of Section 16.2

  42. Colligative Properties of Solutions The physical properties of a solution differ from those of the pure solvent used to make the solution. Some of these differences in properties have little to do with the specific identity of the solute. They depend upon the number of solute particles in the solution. Colligative Property – a property that depends only upon the number of solute particles, and not upon their identity.

  43. Colligative Properties of Solutions • Three important colligative properties of solutions are: • Vapor pressure lowering • Boiling point elevation • Freezing point depression • Vapor pressure is the pressure exerted by a vapor that is in dynamic equilibrium with its liquid in a closed system. • A solution that contains a solute that is nonvolatile (not easily vaporized) always has a lower vapor pressure than the pure solvent.

  44. Colligative Properties of Solutions In a pure solvent, equilibrium is established between the liquid and the vapor. In a solution, solute particles reduce the number of free solvent particles able to escape the liquid. Equilibrium is established at a lower vapor pressure. Ionic solutes that dissociate, such as sodium chloride and calcium chloride, have greater effects on the vapor pressure than does a non-dissociating solute such as glucose. Each formula unit of CaCl2 produces three particles in solution, a calcium ion and two chloride ions.

  45. Colligative Properties of Solutions The decrease in a solution’s vapor pressure is proportional to the number of particles the solute makes in solution. 3 moles of NaCl dissolved in H2O produce 6 mol of particles - each formula unit dissociates into 2 ions 3 moles of CaCl2 dissolved in H2O produce 9 mol of particles - each formula unit dissociated into 3 ions 3 moles of glucose dissolved in water produce 3 mol of particles – glucose does not dissociate.

  46. Colligative Properties of Solutions The vapor pressure lowering caused by 0.1 mol of NaCl in 1000 g of water is twice that caused by 0.1 mol of glucose in the same quantity of water. The vapor pressure lowering caused by 0.1 mol of CaCl2 in 1000 g of water is three times that caused by 0.1 mol of glucose in the same quantity of water. The decrease in a solution’s vapor pressure is proportional to the number of particles the solute

  47. Freezing-Point Depression When a substance freezes, the particles of the solid take on an orderly pattern. The presence of a solute in water disrupts the formation of this pattern because of the shells of water of solvation. (water molecules surround the ions of the solute) As a result, more KE must be withdrawn from a solution than from the pure solvent to cause the solution to solidify. The freezing point of a solution is lower than the freezing point of the pure solvent.

  48. Freezing-Point Depression Freezing-Point Depression – the difference in temperature between the freezing point of a solution and the freezing point of the pure solvent. Freezing-point depression is another colligative property. The magnitude of the freezing-point depression is proportional to the number of solute particles dissolved in the solvent and does not depend upon their identity. The addition of 1 mol of solute particles to 1000 g of water lowers the freezing point by 1.86ºC.

  49. Freezing-Point Depression If you add 1 mole (180g) of glucose to 1000 g of water, the solution freezes at -1.86ºC. If you add 1 mol (58.5g) of NaCl to 1000 g of water, the solution freezes at -3.72ºC, double the change for glucose. This is because 1 mol NaCl produces 2 mol particles and doubles the freezing point depression. Salting icy surfaces forms a solution with the melted ice that has a lower freezing point than water. (antifreeze also)

  50. Reminders Ionic compounds and certain molecular compounds, such as HCl, produce two or more particles when they dissolve in water. Most molecular compounds, such as glucose, do not dissociate when they dissolve in water. Colligative properties do not depend on the kind of particles, but on their concentration. Which produces a greater change in colligative properties – an ionic solid or a molecular solid? An ionic solid produces a greater change because it will produce 2 or more mole of ions for every mol of solid that dissolves.

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