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Solutions

Solutions. We all want them – but few of us have them. Solutions. A solution is a completely homogeneous mixture. Solutions have one or more solutes (dissolvee) and a solvent (dissolver). When a substance dissolves it is soluble

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Solutions

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  1. Solutions We all want them – but few of us have them.

  2. Solutions • A solution is a completely homogeneous mixture. • Solutions have one or more solutes (dissolvee) and a solvent (dissolver). • When a substance dissolves it is soluble • Liquids that mix are miscible, those that don't are immiscible

  3. Characteristics of solutions • Fluid solutions are clear (transparent). • Dissolving is a physical change. Dissolved substances retain their chemical identities. • The physical attributes of the solvent and solute will be altered (physical state, mp, etc).

  4. Types of solutions • Solutions can involve any state of matter. • Liquid/liquid • Solid/solid • Gas/gas • Solid in liquid • Gas in liquid

  5. Aqueous solutions - ionic • Solvation of ionic substances

  6. Aqueous solutions - ionic • Solvation involves dissociation of the solid ionic solute.

  7. Aqueous solutions - ionic • The energy needed to break strong ionic bonds comes from the energy liberated when solvent/solute interactions are formed. • Water, being polar, has partially positive and negative regions which are attracted to the positive and negative ions of the ionic solute.

  8. Aqueous solutions - ionic • Ionic salts are insoluble in non-polar solvents

  9. Aqueous solutions - molecular • Solvation of molecular substances. • Molecular substances do not have strong interactions to overcome. • Dissociation in solution does not occur. • Neutral molecules are separated from each other. • Molecules with polar regions will be soluble in water. • Examples – ethanol (CH3CH2OH), sucrose

  10. Aqueous solutions - molecular • Sucrose

  11. Aqueous solutions - molecular • Polar substances are soluble in water because the polar water molecule can attach to the polar portions of the solute molecule. • Nonpolar molecules are not soluble in water because water is more attracted to itself than the potential solute. • Nonpolar molecules usually do not contain oxygen – example methane (CH4).

  12. Aqueous solutions - molecular • Molecule polarity depends on bond polarity and geometry. • Bond polarity depends on electronegativity differences. • Miscibility – Liquid solutes • Liquids which mix in all proportions are said to be miscible. • The liquid present in larger amounts is considered the solvent. • When mixing unlike liquids, volumes do not necessarily add.

  13. Solubility • Rules of Thumb • Like dissolves like • Polar dissolves polar, nonpolar dissolves nonpolar • Measuring solubilty • Accepted unit is g/100 g solvent. • Temperature must be specified – may be 0º C, may be 20º C, may be 100ºC

  14. Solubility graph

  15. Concentration • Per cent concentration • Weight % (w/w) • w/w% = (mass solute)/(mass solution)x100% • mass solution = mass solute + mass solvent • Volume % (v/v) • v/v% = (volume solute)/(volume solution)x100% • volume solution ≠ volume solute + volume solvent!!

  16. % concentration problem • A solution is prepared by mixing 1.00 g ethanol (C2H5OH) with 100.0 g water to give a final volume of 101 mL. Calculate the w/w and v/v % concentrations. (The density of ethanol is 0.789g/mL)

  17. Solution to concentration problem • w/w: [1.00g eth/101g soln]x100% = 0.990%w/w • v/v: volume of ethanol is 1.00g/0.789g/mL = 1.27mL [1.27mL/101mL]x100% = 1.26%v/v

  18. Molarity • Unit: moles/(liter solution) • Symbol: M • Making a solution: Solute is measured, dissolved in a small amount of water and diluted to desired volume in a volumetric flask.

  19. Molarity examples • What mass of KCl is necessary to make 300 mL of a 0.15 M solution? • Solution: 0.15 M x 0.300 L = 0.045 mol KCl molar mass of KCl = 74.6 g/mol 0.0450 mol x 74.6 g/mol = 3.4 g KCl

  20. Molarity example • You have 458 mL of a 0.29 M solution of sodium hydroxide. What mass of sodium hydroxide is contained therein? • Solution: 0.458 L x 0.29 M = 0.13 mol NaOH 0.13 mol x 40. g/mol = 5.2 g NaOH

  21. Molality • Unit: (moles solute)/(kg solvent) • Symbol: m • Examples: How would you prepare a 0.17 m solution of sodium phosphate using 800. mL water? • Solution: 0.17 mol/kg x .800 kg = 0.14 mol Molar mass of Na3PO4 =163.9 g/m 0.14 mol x 163.9 g/mol = 23g Na3PO4

  22. Molality example • How would you make 1200 g of a 0.235 m calcium chloride solution? • Solution: CaCl2: 111 g/mol • 1.200 kg = mass solvent + mass solute • mass solute = x • 0.235mol CaCl2 =x/111g/mol 1 kg solvent (1.200 kg – x) • Multiply both sides by 111 (111) (111)

  23. Molality example solution • 0.0261kg CaCl2 = x_____ 1 kg solvent (1.200 kg – x) (Note that the mass of the solute must be in kg) • .0261(1.200 – mass solute) = mass solute • .03132–.0261(mass solute) = mass solute • .03132/1.0261 = mass solute • = 0.0305 kg CaCl2 (30.5g)

  24. Mole Fraction • XA = nA/(nA + nB) • No unit (all units cancel) • Example. What is the mole fraction of HCl in concentrated hydrochloric acid (37%w/w)? 37% = 37g HCl/(37g HCl + 63g H2O) • Change all masses to moles (37/36.5)/(37/36.5 + 63/18) • XHCl = 0.22 (22% of the particles in concentrated HCl are hydrogen chloride, neglecting dissociation)

  25. Factors affecting solubility • Temperature • Increase in temperature usually means increased solubility for a solid or liquid solute. • Supersaturation • A solution carefully cooled below temperature where solute is soluble • Disturbing solution or adding seed crystal causes rapid crystallization with attendant evolution or absorption of heat • Gases become less soluble with increasing temperature.

  26. Factors affecting solubility • Nature of solute and solvent • Relative polarities (“like dissolves like”) • Temperature – higher temperature means higher solubility (solids) or lower solubility (gases) • Factors affecting rate of solution • Temperature – higher temperature means faster dissolving • Particle size – smaller particles mean faster solution • Stirring – stirring increases solution rate

  27. Factors affecting solubility • Pressure • Pressure does not affect the solubility of solid or liquid solutes. • Gases become more soluble with increasing pressure • Henry's Law: Gas solubility is directly related to pressure P = kC P = pressure of dissolved gas over solution k= constant characteristic of system C = concentration (solubility)

  28. Henry’s Law problem • A soft drink is bottled so that a bottle at 25ºC contains CO2 at a pressure of 5.0 atm over the liquid. Assuming that the partial pressure of CO2 in the atmosphere is 4.0x10-4atm, calculate the equilibrium concentration of CO2 in the soda both before and after the bottle is opened. The Henry's Law constant for CO2 in aqueous solution is 32 Latm/mol at 25ºC.

  29. Solution to Henry’s Law problem • Before opening: P = kC; k = 32 Latm/mol, and P = 5.0 atm. • So CCO2 = P/k = 5.0atm/32 Latm/mol = 0.16M • After opening: P = kC; k = 32 Latm/mol, and P = 4.0x10-4 atm. • So CCO2 = P/k = 4.0x10-4atm/32 Latm/mol = 1.3x10-5M

  30. Reactions in solution • Precipitates and ionic equations • Overall equation Pb(NO3)2(aq)+2KI(aq)PbI2(s)+2KNO3(aq) • Ionic equation Pb+2(aq)+2NO3‑(aq)+2K++2I‑(aq)PbI2(s)+2K++2NO3‑(aq) • Net ionic equation – eliminate spectator ions Pb+2(aq) + 2I‑(aq) PbI2(s)

  31. Reactions in solution • Other reactions may be driven by the formation of a gas. 2HCl(aq) + Na2CO3(aq) CO2(g) + H2O(l) + 2NaCl(aq) overall equation 2H++2Cl‑(aq)+2Na++CO3‑2(aq)CO2(g)+H2O(l)+2Na+ +2Cl‑(aq) ionic equation 2H+ + CO3‑2(aq)  CO2(g) + H2O(l) net ionic equation

  32. Reactions in solution • Write overall, ionic and net ionic equations for the following reaction. zinc metal reacts with hydrochloric acid to give aqueous zinc chloride and hydrogen gas Zn + 2HCl(aq)  ZnCl2 + H2 Zn + 2H+(aq) + 2Cl-(aq) Zn+2(aq) + 2Cl-(aq) + H2(g) Zn + 2H+(aq) Zn+2(aq) + H2(g)

  33. Reactions in solution • Reactions can also be driven by formation of a complex, usually brightly colored. Cu2+(aq ) + 4 NH3(aq ) Cu(NH3)42+(aq ) 0.1M Cu(II) 0.1M Cu(II) with ammonia added

  34. Reactions in solution • Iron solutions form a complex with thiocyanate (SCN-) Fe+3 + SCN- FeSCN+2 0.1M Fe+3 with thiocyanate added 0.1M Fe+3

  35. Reactions in solution • Formation of a small covalent molecule will also drive a reaction. 2KOH(aq) +H2SO4(aq) K2SO4(aq) +H2O • Use the solubility table (Appendix D) to decide how to write reaction equations for precipitation reactions. • K2S(aq) + Pb(NO3)2(aq) • NaOH(aq) + CuSO4(aq) PbS(s) + 2KNO3(aq) Na2SO4(aq)+Cu(OH)2(s)

  36. Colligative properties • From Latin “colligare”, to bind together • Colligative properties depend on the number of solute particles present • Molecular substances give one mole of particles per mole substance • Salts give more than one mole particles per mole substance because of dissociation • Number of particles per formula unit = i (van’t Hoff factor)

  37. Colligative properties • Vapor pressure lowering • Solute particles take the place of some of the solvent particles at the surface • Vapor pressure of liquid is lowered by presence of a solute • Extent of lowering depends on number of solute particles present

  38. Colligative properties • Boiling point elevation • Solutions have higher boiling points than pure solvents. This is true with solid solutes and heavier liquid solutes. • Other liquid solutes may form azeotropes, which are mixtures with lower boiling points than either solute or solvent – example 95% ethanol/water.) • Solutes raise the boiling point of liquids because they lower the vapor pressure. When Pv = Pa, vaporization (boiling) occurs.

  39. Colligative properties • Boiling point elevation depends on the number of particles present. (van't Hoff factor, i) • Sodium chloride elevates the bp of water twice as much per mole as sucrose, for it makes two particles per mole. (i = 2 for dilute solutions) NaCl Na+ + Cl-

  40. Colligative properties • Freezing point depression • Solutes lower the freezing point of liquid solvents. Solutes interfere with solvents’ ability to associate as a solid and crystallize. • Freezing point depression also depends on number of particles. • Osmosis and osmotic pressure • Solutions of different concentrations on either side of a semi-permeable membrane exert a net pressure toward the more concentrated solution

  41. Colligative properties • Solvent moves so as to dilute the more concentrated solution. Osmotic pressure is the pressure exerted by the greater height of solution on the concentrated side.

  42. Colligative properties • Reverse osmosis is used to purify water.

  43. Colligative properties • Calculations • Colligative property calculations use molality as the unit of concentration, for it does not depend on volume. • Colligative properties vary directly with concentration. • BP elevation: • Tb = ikbm, where kb = the boiling point elevation constant for that liquid.

  44. Colligative properties • kb for water is 0.51ºC/molal • FP depression: • Tf = ikfm • kf for water is 1.86ºC/molal • Compare the freezing points of 1.5 molal aqueous solutions of NaCl and CaCl2. • NaCl: Tf = (2)1.86ºC/m (1.5m) = 5.6ºC. FP = - 5.6ºC • CaCl2: Tf = (3)1.86ºC/m (1.5m) = 8.4ºC FP = - 8.4ºC

  45. Colligative properties • Molecular weight calculations • Solving for molality can yield the molecular weight of a substance if the mass is known. • Example. 1.235 g of a molecular substance dissolved in 100.0 g benzene lowers the freezing point of benzene by 0.468ºC. What is the molecular weight of the substance? kf for benzene is 5.12 ºC/molal.

  46. Colligative properties • Solution: Tf = ikfm m = mol/Kg solvent; mol = mass/mm Tf = ___ikfmass___ (mm)(Kg solvent) mm = ___ikfmass___ (Tf)(Kg solvent) mm = _(1)(5.12ºC/m)(1.235g)____ = 135g/mol (0.100Kg benzene)(0.468ºC)

  47. Nonhomogeneous mixtures • Suspensions • Particles larger than 103 nm (1 m or 10-6 m) • Particles will settle out on standing • Often opaque • Examples: paint, muddy water, orange juice • Colloids • Particles between 1 nm and 1m • Particles will not settle out on standing

  48. Nonhomogeneous mixtures • May appear clear, but exhibits Tyndall effect • Examples: fog, blood, milk, soapy water

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