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Chapter 8 — Solutions and Their Behavior

Chapter 8 — Solutions and Their Behavior

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Chapter 8 — Solutions and Their Behavior

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  1. Chapter 8 — Solutions and Their Behavior

  2. Some Terminology • Solution • A solution is a homogeneous mixture (not possible to see boundaries between components) that consists of one or more solutes uniformly dispersed at the molecular or ionic level throughout a medium known as the solvent • Solvent is present in the larger amount • Solute is present in smaller amount than the solvent • Examples: Liquids: NaCL (sodium chloride + water) 0.89% 99.11% (solute) (solvent) Gases: Air is a solution of N2, O2, few other minor gases

  3. Units of Concentration 1 • Molality (m) aka “molal concentration” Solvent= does not include solute • Molarity (M) aka “molar concentration” Most common concentration unit in Chemistry Solution=solute+solvent

  4. Example • What is the molarity of a solution prepared by dissolving 45.0 grams of NaCl in enough water to give a total volume of 489 mL? Na = 23 a.m.u. Cl = +35 a.m.u. 58 a.m.u. 1mol NaCl = 58g 45g X 1 mol = 0.7759 mol of NaCL 58g 0.7759 mol of NaCl = 1.6 M 0.489 L

  5. Molality vs. Molarity • Molality is never equal to molarity • But the difference becomes smaller as solutions become moredilute(denominators are very similar) • Molarity is more useful when dealing with solution stoichiometry • Molality is more appropriate for dealing with physical chemistry • Question: which is temperature-dependent? • Molarity depends on temperature. Molarity will decrease as temperature increases since the amount of the solution will decrease (from evaporation). Temp  M • Molality does not depend on temperature since mass (kg) does not change with temperature

  6. Units of Concentration 2 • Mole Fraction ( ) • Percent by Volume (% w/v) AND Percent by Weight (%w/w)

  7. Example • Calculate the percent by weight NaCl in a solution comprised of 45.0 g NaCl and 457 g of water. % (w/w) = grams of solute x 100% = 45g x 100% = 8.96% grams of solution 502g Solute = 45g of NaCl Solution = 457g of NaCl (solute) + H2O (solvent) = 502g

  8. Units of Concentration 3 • Parts per million (ppm): Extremely dilute solutions. Compares amount of solute to a million parts of solution (rather than 100 parts). • Parts per billion (ppb) Even more extremely dilute solutions. Compares amount of solute to a billion parts of solution (rather than 1million parts).

  9. Converting Between Units • Every concentration unit is a ratio of two quantities • Pick a sample size This fixes one of the two quantities • Use the factor-label method (dimensional analysis a.k.a. “conversion factor”) to systematically convert the given quantities into the desired quantities

  10. Some More Terminology • Solubility • Solubility is the amount of solute that will dissolve in a given amount of solvent at a given temperature • Saturated • A saturated solution contains the maximum amount of a solute, as defined by its solubility (no more solute will dissolve in a solution that is already saturated with that solute) • Supersaturated • A solution contains more solute than allowed by the solubility. Not a stable system since there is more solute dissolved than the solvent can “accommodate” and the excess solute will come out of solution and will crystallize as a solid (e.g., calcium oxalate or CaPO4 will make kidney stone) or it will separate as a liquid or it will bubble out as a gas (e.g., soda)

  11. Some More Terminology • Miscible • Two liquids are miscible if they are soluble in each other in all proportions • Question: Can a solution be both saturated and dilute? Yes, just remove some of the solvent. If you want to see this for yourself, mix a little salt and water together. Then leave it stand so most of the water evaporates. You have saturated a dilute solution.

  12. Solubility Guidelines • Like dissolves like • Polar solutes are more soluble in polar solvents • Nonpolar solutes are more soluble in nonpolar solvents

  13. Heat of Solution (Energy change or Enthalpy of Solution) ΔH°solution is defined as the enthalpy (energy) change that accompanies dissolving exactly 1 mole of solute in a given solvent Enthalpy = Heat (as long as pressure remains constant) • Some substances have positive (endothermic) ΔH°solution and some have negative (exothermic) ΔH°solution • Endothermic: energy flows INTO the system (gain of energy, temperature of the system decreases) • Exothermic: energy flows OUT of the system (loss of energy, temperature of system increases) ΔH°solution is the sum of two terms: Lattice Energy and Solvation Energy

  14. LE and Solvation • Lattice energy is the energy released when molecules or ions settle into a crystalline lattice. It is inherently exothermic (opposite charges coming towards each other) • The opposite: Pulling all the ions away from each other in an ionic substance in the solid state requires overcoming the lattice energy (endothermic) • Solvation (hydration) energy is the energy released when an ion (or molecules) settles into a sphere of solvent molecules • Solvation is inherently exothermic (opposite charges coming towards each other also) • Both LE and SE result from IM forces

  15. Sodium Chloride: Exothermic DH°solution The magnitude of the lattice energy is lower than the magnitude of the solvation energy, in going from solid sodium chloride to a solution of sodium chloride so the energy of this system has decreased (it has undergone an exothermic change)

  16. Ammonium Chloride: Endothermic DH°solution The magnitude of the lattice energy is greater than the magnitude of the solvation energy, in going from solid ammonium chloride to a solution of ammonium chloride so the energy of this system has increased (it has undergone an endothermic change)

  17. Effect of Pressure on Solubility • Gaseous solute • As P increases, solubility increases • Henry’s Law: S = kHPgas S=solubility, k=Henry’s constant, P=partial pressure of the gas • Liquid and Solid Solutes • P has negligible effect (liquids and solids are not very compressible)

  18. Example • The Henry’s law constant for oxygen in water is 0.042g/L/atm at 25°C. What is the solubility (in grams per mL) of O2 in pure water when the atmospheric pressure is 740 torr. In air, the mole fraction of oxygen is 0.21. Ptotal = 740 torr x 1 atm = 0.974 atm 760 torr PO2 = 0.21 x 0.974 atm = 0.204 atm S = 0.042g/L/atm x 0.204 atm = 0.0086 g/L = 8.6 mg/L

  19. Colligative Properties of Solutions • A colligative property depends only on the number of solute particles, not the identity of the solute particles • Colligative properties include: • Vapor pressure of a solution decreases with increasing [solute] • Boiling point of a solution increases with increasing [solute] • Freezing point of a solution decreased with increasing [solute] • Osmotic pressure of a solution increases with increasing [solute]

  20. Vapor Pressure Decreases • Introduce a solute (solute will “plug escape sites” of molecules that are at the surface of the solution “ready to jump out” and become gas)

  21. Raoult’s Law (vapor pressure of a volatile component of a solution (P) is equal to the vapor pressure of the pure substance Po) times the mole fraction (X) of that substance) • The vapor pressure of a solution is given by Raoult’s Law

  22. Example (Textbook page 206) • What is the vapor pressure of a 5% (by mass) sugar solution at 45°C? (P°water = 71.9 torr)

  23. Example (Textbook, page 206) • What is the total vapor pressure of a solution comprised of 0.75 moles of benzene and 0.45 moles hexane at 60°C? (P°benzene= 390 torr; P°hexane= 580 torr)

  24. Boiling Point Elevation(BP: Temperature at which the vapor pressure of the material is equal to the ambient pressure) • As vapor pressure goes down, boiling point goes up ΔTbp = Kbpmsolute ΔTbp is the boiling point elevation Kbp is the boiling point (ebullioscopic) constant msolute is the molality of all solute particles

  25. Freezing Point Depression(Freezing point: temperature at which the liquid phase of the material is in equilibrium with the solid phase (aka melting point) ΔTfp = Kfpmsolute is ΔTfp the freezing point depression Kfp is the freezing point (cryoscopic) constant msolute is the molality of all solute particles

  26. Ionic Solutes • When ionic solutes dissolve, they dissociate into solvated ions • Each ion counts as a particle for colligative properties

  27. Osmosis • Osmosis is diffusion of water through a semipermeable membrane • Solute particles are too big (or too polar) to make it across the membrane • This is how water gets moved around cells

  28. Tonicity • Isotonic solutions have equal concentrations of solute particles • A hypertonic solution has a greater concentration of solute • A hypotonic solution has lower concentration of solute

  29. Example Hypertonic solution Less water Hypotonic solution More water

  30. Osmotic Pressure (Increases with increasing solute concentration) • Osmotic pressure (P) results from the potential drive for the concentration of water to equalize P V = nRT Or P = MRT A 1.0 M solution of glucose exerts an osmotic pressure of 22.4 atm at 25°C

  31. Question • What will happen to a red blood cell when it is placed into pure water? Cells are isotonic with normal saline (0.89% NaCl)

  32. Question • What will happen to a red blood cell when it is placed into 10% aqueous sodium chloride? Cells are isotonic with normal saline (0.89% NaCl)

  33. Colloids • Colloids are not true solutions • Particle size is on the order of 10 to 200 nm • Might be super-sized molecules (e.g., proteins) or aggregates of ions • Colloidal particles cannot be filtered and do not settle out of solution • Colloids exhibit the Tyndall effect (whereas solutions don’t) • The particles in a colloid are large enough to scatter light passing through) • Examples: • blood, milk, jelly, propofol

  34. Surfactants • A surface active agent breaks the surface tension of water • Surfactants improve a solvent’s ability to be a solvent • Soaps and detergents are common surfactants

  35. polar head greasy tail Soaps and Detergents • A soap is prepared by hydrolyzing fat with alkali • A detergent is a synthetic chemical with a structure similar to soap • Both have a polar (hydrophilic) head and a non-polar (hydrophobic, greasy) tail

  36. Monolayers • The greasy tails stick out of the surface of water • This breaks down the surface tension of water

  37. Bilayers • The tails can dissolve in each other in a bilayer • This structure is used in cell membranes

  38. Micelles • The tails can dissolve in each other forming a sphere • This creates a non-polar microenvironment in the water

  39. Thank you!