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Chapter 13 Solutions

Chapter 13 Solutions. Types of Mixtures. Heterogeneous and Homogeneous Mixtures. Solutions. Soluble – capable of being dissolved

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Chapter 13 Solutions

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  1. Chapter 13Solutions

  2. Types of Mixtures • Heterogeneous and Homogeneous Mixtures

  3. Solutions • Soluble – capable of being dissolved • Solution – a homogeneous mixture of two or more substances in a single phase; a “well-mixed” mixture; can be a combination of any phases of matter. Solute particles are less than 1 nm in diameter. • Examples – salt water, air, brass • Solvent – the dissolving medium in the solution • Ex. Water is the solvent in salt water. • Solute – the substance dissolved in the solution • Ex. Salt is the solute in salt water.

  4. Solute-Solvent Combinations • Ex.

  5. Suspensions • A mixture that contains particles in a solvent that are so large that they settle out unless the mixture is constantly stirred or agitated. • Particles are generally greater than 1000 nm in diameter. • Suspensions may reflect or scatter light. • Not transparent, and is hetergeneous

  6. Colloids • Particles that are intermediate in size between those in solutions and suspensions form mixtures called colloidal dispersions, or colloids. • Do not separate upon standing and are heterogeneous • Particles size between 1 nm and 1000 nm in diameter • Colloidal particles make up the “dispersed phase”, while water is the “dispersing medium” • Term colloid comes from Greek word, “kola” which means “glue”. Named by Thomas Graham in 1861; founder of colloid chemistry and dialysis (medical treatment).

  7. Classes of Colloids

  8. Distinguishing Characteristics of Colloids • Tyndall Effect - Light is scattered by colloidal particles dispersed in a transparent medium. • Ex. Headlights shine through foggy weather at night. • Can be used to distinguish a colloid from a solution when the colloidal particles are otherwise too small to see.

  9. Distinguishing Characteristics of Colloids • Brownian motion – flashes of light or scintillations are observed under a microscope; is a collision of the molecules of the dispersion medium with small dispersed colloidal particles. These conditions prevent settling. • Coagulation – colloidal particles are charged by absorbing ions from the dispersing medium onto their surface.

  10. Solutes • Electrolytes vs. Nonelectrolytes – classified according to whether they yield molecules or ions in solution. • When an ionic compound dissolves, the positive & negative ions separate from each other and are surrounded by water molecules. • Electrolyte- a solute that dissolves in water to produce ions in solution that can conduct electricity. • Example – NaCl is an electrolyte • Nonelectrolyte – a solute that dissolves in water to produce a solution that does not conduct electricity. • Example – sugar is a nonelectrolyte

  11. Solution Process • Creating a solution can be helped by 3 actions: • Increasing the surface area exposed to solvent or powdering • Stirring or agitating • Heating the solvent

  12. Solubility Equilibrium • The physical state in which the opposing processes of dissolution and crystallization of a solute occur at equal rates.

  13. Saturated vs. Unsaturated • A solution that contains the maximum amount of dissolved solute is described as a saturated solution. • A saturated solution in a closed system is at equilibrium • A solution that contains less solute than a saturated solution under the existing conditions is an unsaturated solution.

  14. Supersaturated Solutions • A solution that contains more dissolved solute than a saturated solution contains under the same conditions. • Will usually become saturated at a lower temperature.

  15. Solubility • Amount of substance to form a saturated solution with a specific amount of solvent at a specific temperature. • Solubility charts generally have grams solute / 100g solvent at a given temperature. • Example: The solubility of sugar is 204g per 100.g of water at 200C.

  16. “Like Dissolves Like” • Polar and ionic solutes like NaCl,LiCl, and MgSO4, tend to dissolve in polar solvents like H2O • Nonpolar solutes tend to dissolve in nonpolar solvents, like gasoline, C8H18, dissolving in benzene, C6H6.

  17. Immiscibility vs. Miscibility • Liquid solutes and solvents that are not soluble in each other are immiscible.

  18. Hydration • Terms – • Hydration – any solution process with water as the solvent.

  19. Hydrates • When crystallized from aqueous solutions, some ionic substances form crystals that incorporate water molecules. • Ex. CuSO4. 5H2O

  20. Effects of Pressure on Solubility • Changes in pressure have very little effect on the solubilities of liquids or solids in liquid solvents. • BUT, increases in pressure increase gas solubilities in liquids. • Gas + solvent solution • Henry’s Law – named after the English chemist William Henry states: The solubility of a gas in a liquid is directly proportional to the partial pressure of that gas on the surface of the liquid.

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

  22. Effects of Temperature on Solubility • Increasing the temperature usually decreases gas solubility, because as the avg KE of the particles in solution increases, the number of solute particles that can escape the attractions of the solvent increases and the gas can escape to return to gas phase.

  23. Temperature vs. Solubility Whether an increase in temperature increases or decreases the solubility of a solid is dependent upon the identity of the solid.

  24. Heat of Solution • The formation of a solution involves an energy change. Sometimes the change is endothermic, sometimes exothermic. • Formation of solutions can absorb heat (endothermic) • Example: potassium iodide in water • Formation of solutions can release heat (exothermic) • Example: lithium chloride in water • The net amount of heat energy absorbed or released when a specific amount of solute dissolves in a solvent is the heat of solution.

  25. Heat of Solution

  26. 2 Methods for Expressing the Concentration of a Solution • The concentration of a solution is a measure of the amount of solute in a given amount of solvent or solution. • 1). Molarity – M; the number of moles of solute in one liter of solution. • Must determine the molar mass of the solute in order to find molarity of a solution • Ex. A 1-molar solution of NaOH contains 1 mole of NaOH in a Liter of solution. (written as 1 M NaOH) • Molarity(M) = amount of solute (mol) Volume of solution (L)

  27. Sample Problem 1 • You have 3.50 L of solution that contains 90.0g of sodium chloride, NaCl. What is the molarity of that solution? • Recall: Molarity (M) = amount of solute (mol) V solution (L) 90.0 g NaCl. 1 molNaCL = 1.54 mol 58.5 g NaCl M = 1.54 mols/3.50 L = .440 M NaCl

  28. Sample Problem 2 • You have 0.8 L of a 0.5 M HCl solution. How many moles of HCl does this solution contain? • Given – volume = 0.8 L concentration of solution = 0.5 M HCl 0.5 M HCl = mols of HCl= 0.4 molHCl 0.8 L

  29. Sample Problem 3 • To produce 40.0g of silver chromate, you will need at least 23.4g of potassium chromate in solution as a reactant. All you have on hand in the stock room is 5 L of a 6.0 M K2CrO4 solution. What volume of the solution is needed to give you the 23.4g K2CrO4 needed for the reaction? • Given: volume of solution = 5 L • Concentration of solution = 6.0 M K2CrO4 • Mass of solute = 23.4g K2CrO4 • Mass of product = 40.0g Ag2CrO4 • Unknown: volume of K2CrO4 solution in L • 23.4g K2CrO4. 1 mol K2CrO4 = .120 mol K2CrO4 194.2 g K2CrO4 • 6.0 M K2CrO4 = .120 mol K2CrO4 = .020 L K2CrO4 solution Volume of solution

  30. Molality • The concentration of a solution expressed in moles of solute per kilogram of solvent. • Molality = m = moles solute mass of solvent (kg) • Ex. A solution that contains 1 mol of solute, NaOH, dissolved in exactly 1 kg of solvent is a “one-molal” solution. (written as 1 mNaOH) • Concentrations are expressed as molalities when studying properties of solutions related to vapor pressure and temperature changes. Molality is used because it does not change with changes in temperature.

  31. Sample Problem 1 • A solution was prepared by dissolving 17.1g of sucrose (table sugar, C12H22O11) in 125g of water. Find the molal concentration of this solution. • Given: solute mass = 17.1g sucrose; • solvent mass = 125g water • Molality = moles solute/mass of solvent(kg) • 17.1g C12H22O11/342 g C12H22O11 = 0.050 mol C12H22O • 0.050 mol sucrose / .125kg water = 0.400 m C12H22O11

  32. Sample Problem 2 • A solution of iodine, I2, in carbon tetrachloride, CCl4, is used when iodine is needed for certain chemical tests. How much iodine must be added to prepare a 0.480 m solution of iodine in CCl4 if 100.0g of CCl4 is used? • Given: 0.0480 m I2 solution; mass of solvent= 100.0g CCl4 • Unknown: mols of I2 • Molality = mols solute/mass of solvent (kg) • 0.480 m I2 x .1000kg CCl4 = .048 mol I2 • .048 mol I2 x 253.8g I2 = 12.2 grams I2

  33. Mass Percentage Concentration • As a percentage value, measures the number of grams of solute in 100 g of solution. • Is a convenient unit when solute is a solid. • May utilize density and volume of a solvent. • Equation: mass percentage (%) = mass of solute mass of solution • = mass of solute mass of solute + mass of solvent

  34. Example 1 • How many grams of K2SO4 would you need to prepare a 5% solution if 1500 g of water are used as the solvent?

  35. Example 2

  36. Ions in Solution • Main Idea: • Ionic Compounds dissociate in aqueous solutions. • Colligative properties of a solution depend only on the number of solute particles present. • The separation of ions that occurs when an ionic compound dissolves is called dissociation.

  37. Moles of Ions • Examples:

  38. Sample Problem • Write the equation for the dissolution of aluminum sulfate, Al2(SO4)3, in water. • How many moles of aluminum ions and sulfate ions are produced by dissolving 1 mol of aluminum sulfate? • What is the total number of moles of ions produced by dissolving 1 mol of aluminum sulfate?

  39. Sample Problem 2 • Write the equation for the dissolution of each of the following in water, and then determine the number of moles of each ion produced as well as the total number of moles of ions produced. • 1 mol ammonium chloride • 1 mol sodium sulfide • 0.5 mol barium nitrate

  40. Precipitation Reactions • No ionic compound is completely insoluble, but compounds of very low solubility can be considered insoluble for most practical purposes.

  41. General Solubility Guidelines

  42. Example • Write a net ionic equation for the double-displacement reaction between (NH4)2S and Cd(NO3)2.

  43. Predicting a precipitate? • 2 dissociation equations are: • (NH4)S(s)  2NH4+ (aq) + S2-(aq) • Cd(NO3)2(s)  Cd2+ (aq) + 2NO3-(aq) • 2 possible products • (NH4)2S(aq) + Cd(NO3)2(aq)  NH4NO3(?) + CdS(?)

  44. (NH4)2S(aq) + Cd(NO3)2(aq)  NH4NO3(aq) + CdS(s)

  45. Net Ionic Equations • Includes only those compounds and ions that undergo a chemical change in a reaction in an aqueous solution. • Steps for writing net ionic equations • 1. convert the chemical equation into a complete ionic equation. • 2. All soluble ionic compounds are shown as dissociated ions in solution. • 3. The precipitates are shown as solids. • 4. The spectator ions are canceled on both sides of the equation.

  46. Net Ionic Equation • Includes only those compounds and ions that undergo a chemical change in a reaction in an aqueous solution. • Complete: • Cd2+(aq) + 2NO3-(aq) + 2NH4+ (aq) + S2-(aq)  • CdS(s) + 2NO3-(aq) + 2NH4+ (aq) • Notice: Ammonium ions and nitrate ions are spectator ions, only. (do not take part in the chem. reaction, and are found in solution.) • Spectator ions are canceled on both sides. • Net ionic equation is: Cd2+(aq) + S2-(aq)  CdS (s)

  47. Ionization • Ions are formed from solute molecules by the action of a polar solvent in a process called ionization. (this creates ions where there were none before) • The extent to which a solute within the solution ionizes depends on the strength of the bonds between the molecules of the solute, and the strength of the attraction between the solute and solvent molecules.

  48. Hydronium ion • H3O+ ion that forms when a covalently bonded molecule containing a hydrogen bonds with a water molecule in aqueous solution.

  49. Electrolytes • .

  50. Strong vs. Weak

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