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Properties of Solutions

This chapter explores the properties of solutions, including how they form, ways of expressing concentrations, and energy changes. It also emphasizes the importance of understanding the difference between dissolution and chemical reactions in solutions.

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Properties of Solutions

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  1. Chapter 11 - Properties of SolutionsProblems: 27, 29, 31, 33, 36, 40, 43, 47-58, 59, 61, 63, 65, 67, 69

  2. Solutions • Solutions are homogeneous mixtures of two or more pure substances. • In a solution, the solute is dispersed uniformly throughout the solvent.

  3. Solutions The intermolecular forces between solute and solvent particles must be strong enough to compete with those between solute particles and those between solvent particles.  2009, Prentice-Hall, Inc.

  4. How Does a Solution Form? As a solution forms, the solvent pulls solute particles apart and surrounds, or solvates, them.  2009, Prentice-Hall, Inc.

  5. How Does a Solution Form If an ionic salt is soluble in water, it is because the ion-dipole interactions are strong enough to overcome the lattice energy of the salt crystal.  2009, Prentice-Hall, Inc.

  6. Ways of Expressing Concentrations of Solutions

  7. mol of solute L of solution M = Molarity (M) • Since volume is temperature-dependent, molarity can change with temperature.

  8. mass of solute total mass of solution Mass Percentage Mass % of solute =  100

  9. moles of A total moles in solution XA = Mole Fraction (X) • In some applications, one needs the mole fraction of solvent, not solute — make sure you find the quantity you need!

  10. mol of solute kg of solvent m = Molality (m) Since both moles and mass do not change with temperature, molality (unlike molarity) is not temperature-dependent.

  11. Changing Molarity to Molality If we know the density of the solution, we can calculate the molality from the molarity and vice versa.  2009, Prentice-Hall, Inc.

  12. gram mL density = Density To convert from grams to milliliters of a solution use the inverse of the density.

  13. Learning Check A solution of phosphoric acid was made by dissolving 10.0 g H3PO4 in 100.0 mL water. The resulting volume was 104 mL. Given water has a density of 1.00 g/cm3, calculate the following for the solution: • Density • Mole fraction • Molarity • Molality  2009, Prentice-Hall, Inc.

  14. # of equivalents L of solution Normality = Normality Normality depends on the nature of the reaction. For example, in an oxidation-reduction reaction the equivalent is the quantity of the oxidizing or reducing agent that can furnish 1 mol of –e . You will not be ask to calculate a normality on the exam.

  15. grams solute volume of solution % = Mass/ volume percent x 100%

  16. volume solute volume of solution % = Volume/ volume percent x 100% Proof: The “proof” of a solution is twice the percentage. For example a 90 proof solution is 45% alcohol.

  17. mass of A in solution total mass of solution mass of A in solution total mass of solution Parts per Million andParts per Billion ppm = Parts per Million (ppm)  106 Parts per Billion (ppb)  109 ppb =

  18. Energy Changes in Solution • Simply put, three processes affect the energetics of solution: • separation of solute particles, H>0, endothermic • separation of solvent particles, H >0, endothermic • new interactions between solute and solvent, H<0, exothermic  2009, Prentice-Hall, Inc.

  19. Energy Changes in Solution The enthalpy change of the overall process depends on H for each of these steps. Hsoln= H 1+H2+H3  2009, Prentice-Hall, Inc.

  20. Why Do Endothermic Processes Occur? Things do not tend to occur spontaneously, (i.e., without outside intervention) unless the energy of the system is lowered.  2009, Prentice-Hall, Inc.

  21. Why Do Endothermic Processes Occur? Yet we know the in some processes, like the dissolution of NH4NO3 in water, heat is absorbed, not released.  2009, Prentice-Hall, Inc.

  22. Enthalpy Is Only Part of the Picture The reason is that increasing the disorder or randomness (known as entropy) of a system tends to lower the energy of the system.  2009, Prentice-Hall, Inc.

  23. Enthalpy Is Only Part of the Picture So even though enthalpy may increase, the overall energy of the system can still decrease if the system becomes more disordered.  2009, Prentice-Hall, Inc.

  24. Student, Beware! Just because a substance disappears when it comes in contact with a solvent, it doesn’t mean the substance dissolved.  2009, Prentice-Hall, Inc.

  25. Student, Beware! • Dissolution is a physical change — you can get back the original solute by evaporating the solvent. • If you can’t, the substance didn’t dissolve, it reacted.  2009, Prentice-Hall, Inc.

  26. CHECK YOUR KNOWLEDGE • Which of the following has been dissolved and which is a chemical reaction that leads to a solution? • 1) NaCl and water • 2) Ni and HCl  2009, Prentice-Hall, Inc.

  27. CHECK YOUR KNOWLEDGE • Which of the following has been dissolved and which is a chemical reaction that leads to a solution? • 1) physical process of solution • 2) chemical reaction that leads to a solution but the nickel is not dissolved!

  28. CHECK YOUR KNOWLEDGE • Silver chloride, AgCl, is essentially insoluble in water. Would you expect a significant change in the entropy of the system when 10 grams of AgCl is added to 500 mL of H2O?

  29. Types of Solutions • Warm-up: Create a tree chart on the different types of solutions. Due in 15 minutes!

  30. Factors Affecting Solubility • Structure Effects or Solute-Solvent Interactions: • Chemists use the axiom “like dissolves like." • Polar substances tend to dissolve in polar solvents. • Nonpolar substances tend to dissolve in nonpolar solvents.

  31. Factors Affecting Solubility miscible: liquids that mix in all proportions of H2O. immiscible: liquids that don’t mix (water and gasoline)

  32. Factors Affecting Solubility • In general, the solubility of gases in water increases with increasing mass. • Larger molecules have stronger dispersion forces.

  33. Factors Affecting Solubility • The solubility of liquids and solids does not change appreciably with pressure. • The solubility of a gas in a liquid is directly proportional to its pressure.

  34. Henry’s Law Sg = kPg where • Sg is the solubility of the gas, • k is the Henry’s Law constant for that gas in that solvent, and • Pg is the partial pressure of the gas above the liquid.  2009, Prentice-Hall, Inc.

  35. Factors Affecting Solubility TEMPERATURE: Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature.  2009, Prentice-Hall, Inc.

  36. Temperature • The opposite is true of gases. • Carbonated soft drinks are more “bubbly” if stored in the refrigerator. • Warm lakes have less O2 dissolved in them than cool lakes.  2009, Prentice-Hall, Inc.

  37. Colligative Properties • Changes in colligative properties depend only on the number of solute particles present, not on the identity of the solute particles. • Among colligative properties are • Vapor pressure lowering • Boiling point elevation • Freezing point depression • Osmotic pressure

  38. Vapor Pressure Because of solute-solvent intermolecular attraction, higher concentrations of nonvolatile solutes make it harder for solvent to escape to the vapor phase.  2009, Prentice-Hall, Inc.

  39. Vapor Pressure Therefore, the vapor pressure of a solution is lower than that of the pure solvent.  2009, Prentice-Hall, Inc.

  40. Raoult’s Law Psoln = XsolventPsolvent where • XA is the mole fraction of compound A, and • PA is the normal vapor pressure of A at that temperature. NOTE: This is one of those times when you want to make sure you have the vapor pressure of the solvent.

  41. Boiling Point Elevation and Freezing Point Depression Nonvolatile solute-solvent interactions also cause solutions to have higher boiling points and lower freezing points than the pure solvent.

  42. Boiling Point Elevation • The change in boiling point is proportional to the molality of the solution: Tb = Kb  m where Kb is the molal boiling point elevation constant, a property of the solvent. Tb is added to the normal boiling point of the solvent.  2009, Prentice-Hall, Inc.

  43. Freezing Point Depression • The change in freezing point can be found similarly: Tf = Kf  m • Here Kf is the molal freezing point depression constant of the solvent. Tf is subtracted from the normal freezing point of the solvent.  2009, Prentice-Hall, Inc.

  44. Boiling Point Elevation and Freezing Point Depression Note that in both equations, T does not depend on what the solute is, but only on how many particles are dissolved. Tb = Kb  m Tf = Kf  m

  45. Colligative Properties of Electrolytes Since these properties depend on the number of particles dissolved, solutions of electrolytes (which dissociate in solution) should show greater changes than those of nonelectrolytes.  2009, Prentice-Hall, Inc.

  46. Colligative Properties of Electrolytes However, a 1M solution of NaCl does not show twice the change in freezing point that a 1M solution of methanol does.  2009, Prentice-Hall, Inc.

  47. van’t Hoff Factor One mole of NaCl in water does not really give rise to two moles of ions.

  48. van’t Hoff Factor Some Na+ and Cl- reassociate for a short time, so the true concentration of particles is somewhat less than two times the concentration of NaCl.  2009, Prentice-Hall, Inc.

  49. van’t Hoff Factor • Reassociation is more likely at higher concentration. • Therefore, the number of particles present is concentration-dependent.  2009, Prentice-Hall, Inc.

  50. van’t Hoff Factor • We modify the previous equations by multiplying by the van’t Hoff factor, i. Tf = Kf  m  i  2009, Prentice-Hall, Inc.

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