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Solutions

Solutions. Chapter 13. Mixtures. Not pure substances Contain more than one kind of element or compound Homogeneous mixtures The same throughout Heterogeneous mixtures Not the same throughout. Solutions. Homogenous mixtures of two or more substances in a single phase

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Solutions

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  1. Solutions Chapter 13

  2. Mixtures • Not pure substances • Contain more than one kind of element or compound • Homogeneous mixtures • The same throughout • Heterogeneous mixtures • Not the same throughout

  3. Solutions • Homogenous mixtures of two or more substances in a single phase • Cannot be separated by physical means

  4. Soluble • Capable of being dissolved • Some substances are soluble in certain liquids, but not others

  5. Parts of a solution • Solute • The thing that gets dissolved • Solvent • The thing that does the dissolving

  6. Dissolving • Surface molecules leave the crystal and mix with the solvent. • Eventually, the solute molecules become evenly distributed throughout the solvent

  7. Types of solutions • See page 396

  8. Suspensions • Heterogeneous mixture • Particles settle out if the mixture isn’t stirred or shaken • Can be separated with filters

  9. Colloids • Between solutions and suspensions • Particles don’t settle • Look cloudy • Scatter light • See page 398

  10. Tyndall Effect • When light is scattered by colloidal particles dispersed in a transparent medium • Can be used to distinguish solutions and colloids

  11. Electrolytes • Yield ions in solutions • Solutions conduct electricity

  12. Nonelectrolytes • Yield neutral molecules in solution • Solutions do not conduct electricity

  13. Discuss • Classify the following as either a heterogeneous or homogeneous mixture. Explain • Orange juice • Tap water • Salt water • Describe one way to prove that a mixture of sugar and water is a solution and that a mixture of sand and water is not a solution

  14. Dissolving a solid in a liquid • Rate of dissolution affected by: • Surface area of solute • Powdered substances dissolve faster • Agitation • Stirring or shaking • Solvent temperature • Heating the solvent

  15. Solubility • For every combination of solvent with a solid solute at a given temperature, there is a limit to the amount of solute that can be dissolved. • Depends on the natures of the solute and solvent and the temperature

  16. Solubility • The amount of substance required to form a saturated solution with a specific amount of solvent at a specified temperature. • Example: for sugar, the solubility is 204 g per 100. g of water at 20. °C. • For gases, pressure must also be specified

  17. Solution equilibrium • The state in which dissolution and crystallization of a solute occur at equal rates

  18. Saturated solution • Contains the maximum amount of dissolved solute • If more solute is added, it does not dissolve

  19. Unsaturated solution • Contains less solute than a saturated solution • If more solute is added, it will dissolve

  20. Supersaturated solution • Contains more dissolved solute than a saturated solution under the same conditions • Unstable • Formed when a saturated solution cools to a lower temperature without crystallization

  21. Like dissolves like • Similar substances will dissolve in each other • Type of bonding • Polarity or nonpolarity • Intermolecular forces between solvent and solute

  22. Ionic compound in water • The attraction between the polar water molecules and the ions is strong enough to pull the crystal apart. • This is called hydration. • When the ions are surrounded by water molecules they are hydrated. • When any solute particle is surrounded by any solvent molecules, it is solvated.

  23. Hydrates • Compounds that crystallized from aqueous solutions and incorporated water molecules into their crystal structure. • Have specific ratios of water molecules. • We can drive off the water by heating, obtaining the anhydrous form.

  24. Nonpolar solvents • Generally can’t dissolve ionic compounds • Intermolecular attraction is too weak

  25. Immiscible • Liquid solutes and solvents that are not soluble in each other • Oil and water

  26. Miscible • Liquids that dissolve freely in one another in any proportion • Ethanol and water

  27. Effects of pressure • Increased pressure increases the solubility of gases in liquids. • Le Châtelier’s Principle • Gas + solvent ↔solution

  28. Henry’s Law • The solubility of a gas in a liquid is directly proportional to the partial pressure of that gas on the surface of the liquid.

  29. Effervescence • The rapid escape of a gas from a liquid in which it is dissolved

  30. Temperature and solubility • For most solids, they will dissolve more in warm solvents. • For gases, they will dissolve more in cool solvents.

  31. Heat of solution • Net amount of heat energy absorbed or released when a specific amount of solute dissolves in a solvent

  32. Discuss • Why would you expect a packet of sugar to dissolve faster in hot tea than in iced tea? • Explain how you would prepare a saturated solution of sugar in water. • How would you then make it a supersaturated solution? • Explain why ethanol will dissolve in water and carbon tetrachloride will not. • If a warm bottle of soda and a cold bottle of soda are opened, which will effervesce more and why?

  33. Concentration • A measure of the amount of solute in a given amount of solvent or solution

  34. Molarity • The number of moles of solute in one liter of solution. • mol/L • 1 M NaOH • One-molar sodium hydroxide • There is 1 mole of NaOH in each liter of solution

  35. Making 1 molar solutions • Dissolve 1 mole of solute in less than 1 L of solvent. • Then add solvent until the total volume is 1 L.

  36. Example • What is the molarity of a 2.0 L solution that is made from 14.6 g of NaCl? • 0.12 M

  37. You try • What is the molarity of a HCl solution that contains 10.0g of HCl in 250 mL of solution? • 1.10 M

  38. You try • How many moles of NaCl are in 1.25 L of 0.330 M NaCl? • 0.412 mol

  39. Molality • Concentration expressed as moles of solute per kg of solvent. • mol/kg • 1 m NaOH • One-molal sodium hydroxide • There is 1 mol of NaOH for each kg of water

  40. Making1 molal solutions • Place 1 mole of solute in a beaker or flask. • Then add exactly 1 kg of solvent.

  41. Example • What is the molality of a solution composed of 13.0 g NaCl dissolved in 500. g of water? • 0.445 m

  42. You try • How many grams of NaCl are needed to prepare a 1.0 m solution using 250 g of solvent? • 15 g

  43. You try • What volume water must be used to make a 0.245 m solution of NaCl that contains 1.0 mol of this salt? • 4.1 L

  44. Dilutions • You can dilute a solution of a given molarity to a lower molarity using the following equation: • M stands for molarity • V stands for volume of solution

  45. Example • You have a large bottle of 6.0 M HCl. You need 60 mL of 1.0 M HCl. How much 6.0 M HCl should you use to make your dilution?

  46. Example continued • To make the solution, add 10 mL of the 6.0 M HCl to less than 50 mL of water. Then, add water until the total volume is 60 mL.

  47. You try • A dilution was made by adding 20 mL of 5.0 M NaOH solution to 80 mL of water. • What is the molarity of the new solution?

  48. You try continued

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