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Unit 12: Properties of Solutions

CHM 1046 : General Chemistry and Qualitative Analysis. Unit 12: Properties of Solutions. Dr. Jorge L. Alonso Miami-Dade College – Kendall Campus Miami, FL. Textbook Reference : Chapter # 14 Module # 2. Solutions.

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Unit 12: Properties of Solutions

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  1. CHM 1046: General Chemistry and Qualitative Analysis Unit 12:Properties of Solutions Dr. Jorge L. Alonso Miami-Dade College – Kendall Campus Miami, FL • Textbook Reference: • Chapter # 14 • Module # 2

  2. Solutions • Solutions (soln) are homogeneous mixtures of two or more pure substances. • The solvent (solv) is present in greatest abundance. • All other substances are solutes (solu). Volumetric flask {PrepASolu}

  3. Solutions • Solutions are homogeneous mixtures of two or more pure substances. • In a solution, the solute(< 50%) is dispersed uniformly throughout the solvent (> 50%). Homogeneous Heterogeneous

  4. Student, Beware!: solution vs. reaction H2O H2O H2 (g) + MgCl2 (aq) Cu(NO3)2(s) Cu(NO3)2(aq) Mg (s) + 2 HCl (aq) Just because a substance disappears when it comes in contact with a solvent, it doesn’t mean the substance dissolved. • 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.

  5. How Does a Solution of Salts in Water Form? The intermolecular forces between solute and solvent particles must be strong enough to compete with those between solute particles and those between solvent particles. Example: {NaCl + H2O*} NaCl (s) Na+(aq) + Cl- (aq) + H20 dissociation

  6. Solubility in water Elements: mostlyinsolublesolids, liquids & gases. Covalent Compounds: many areinsolublegases, polar covalent are mostly soluble. • Ionic Compounds: many are soluble. • SOLUBILITY RULES: for Ionic Compounds (Salts) • All salts of alkali metals (IA) are soluble. • All NH4+ salts are soluble. • All salts containing the anions: NO3-, ClO3-, ClO4-, (C2H3O2-) are soluble. • All Cl-, Br-, and I- are solubleexcept forAg+, Pb2+, and Hg22+ salts. • All SO42- are solubleexcept forPb2+, Sr2+, and Ba2+. • All O2- are insolubleexcept forIA metalsCa2+, Sr2+, and Ba2+ salts. • {Soluble metal oxides form hydroxides: CaO Ca 2+ + 2OH-} • 7. All OH- are insolubleexcept for IA metals, NH4+ & slightly soluble Ca 2+Ba2+ & Sr2+ • 6. All salts containing the anions: CO32-, PO43-, AsO43-, S2- and SO32- are insolubleexcept fro IA metals and NH4+ salts. • 7. For salts containing the anions not mentionedabove (e.g., CrO42-, Cr2O72-, P3-, C2O42- etc.) assume that they areinsolubleexcept for IA metals and NH4+ salts, unless, otherwise informed. H2O

  7. Energetics of SolutionsHeats (Enthalpy) of Solution (Hsoln) Exothermic solution process Hot & Cold Packs H2O CaCl2 (s) Ca 2+(aq) + 2Cl-(aq) + Heat Hsoln = - 81.3 kJ/mol NH4NO3 +H2O CaCl2 + H2O • Endothermic solution process H2O Heat + NH4NO3 (s) NH4+(aq) + NO3-(aq) Hsoln = + 25.7 kJ/mol

  8. How Does a Solution Form?(3 events) (3) Formation of Ion-dipole interactions (2) Separation of solvent molecules (1) Separation of solute molecules 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. H2O H Lattice Energy +788 kJ/mol Endothermic NaCl Crystal Lattice Hvap +41 kJ/mol Endothermic Exothermic H- bonding Energy of SolutionHsoln = + or -

  9. Energy Changes in Solution {EnerSoln} The enthalpy change of the overall process depends on H for each of these steps. - +

  10. Why do Endothermic Reactions occur Spontaneously? Exothermic solution process Endothermic solution process H2O CaCl2 (s) Ca 2+(aq) + 2Cl-(aq) Hsoln = - 81.3 kJ/mol H2O NH4NO3 (s) NH4+(aq) + NO3-(aq) Hsoln = + 25.7 kJ/mol

  11. Concentration of Solutions (General) • Unsaturated • Less than the maximum amount of solute for that temperature is dissolved in the solvent. {UnsatSoln} • Saturated • Solvent holds as much solute as is possible at that temperature. • Dissolved solute is in dynamic equilibrium with solid solute particles. Solubility of NaCl in H2O = 35.9 g/100 mL (25 °C) NaC2H4O2 = 76 g/100 ml (0°C)

  12. Concentration of Solutions (General) {*SuperSat1} {*SuperSat2} • Supersaturated • Solvent holds more solute than is normally possible at that temperature. • These solutions are unstable; crystallization can usually be stimulated by adding a “seed crystal” or scratching the side of the flask.

  13. Factors Affecting SolubilityA. Solids and Liquids: • Polarity • Temperature B. Gases : • Molar Mass • Pressure – Henry’s Law • Temperature

  14. Factors Affecting Solubility of Solids and Liquids: (1) Polarity • Chemists use the axiom “like dissolves like”: • Polar substances tend to dissolve in polar solvents. • Nonpolar substances tend to dissolve in nonpolar solvents. hydrophobicwater-hating, nonpolar, hydrophilicwater-loving, polar, (hydrocarbon chain) [Know molecular shapes and polar molecules] {Water + Oil}

  15. Factors Affecting Solubility of Solids and Liquids: (1) Polarity • Chemists use the axiom “like dissolves like”: • Polar substances tend to dissolve in polar solvents. • Nonpolar substances tend to dissolve in nonpolar solvents. Polar Non-Polar Polar heads Non-Polar tails [Know molecular shapes and polar molecules]

  16. Factors Affecting Solubility of Solids and Liquids: (1) Polarity Glucose (which has hydrogen bonding) is very soluble in water, while….. …..cyclohexane (which only has dispersion forces) is not.

  17. Factors Affecting Solubility of Solids and Liquids: (1) Polarity • Vitamin A is soluble in nonpolar compounds (like fats). • Vitamin C is soluble in water. {*Iodine + Water then + CCl4 }

  18. Factors Affecting Solubility of Solids and Liquids: (2) Temperature Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature.

  19. Factors Affecting Solubility ofGases in Solution: (1) Molar Mass • In general, the solubility of gases in water increases with increasing molar mass. • Larger molecules have stronger London dispersion forces. Which gas dissolves best in water? g-MM 28 g/n 28 g/n 32 g/n 40 g/n 84 g/n Why?

  20. Factors Affecting Solubility of Gases in Solution: (2) Pressure • The solubility of liquids and solids does not change appreciably with pressure. • The solubility of a gas (Sg) in a liquid is directly proportional to its pressure (Pg). {Henrys Law} SgPg

  21. Factors Affecting Solubility of Gases in Solution: (2) Pressure Henry’s Law: {movie} SgPg where • Sg is the solubility of the gas; • Pg is the partial pressure of the gas above the liquid. • k is the Henry’s law constant for that gas in that solvent; Sg = kPg Pg Sg @ 250C, PN2= 0.75 atm, SN2= 0.00053 M. What is SN2 when PN2 = 1.50 atm? Ans.= 0.00106 M

  22. Factors Affecting Solubility of Gases in Solution: (3) Temperature • The opposite is true of gases: • Gases in carbonated soft drinks are less soluble when warmed. • Warm lakes have less O2 dissolved in them than cool lakes.

  23. Colligative Properties How does the presence of a solute change the properties of a solvent? • Vapor pressure? • Boiling point? • Freezing (melting) point? • Osmotic pressure? Solution: Solvent + Solute Pure Solvent

  24.  =MRT Colligative Properties • Properties that depend only on the CONCENTRATION (number of solute particles present), not on the identity of the solute particles. • Vapor pressure lowering • 2. Boiling point elevation • 3. Freezing point • depression • 4. Osmotic pressure PA =XAPA Tb=Kb  m Tf=Kf m Pure solvent Solvent + Solute

  25. Ways of Expressing Concentrations of Solutions (1) Mass Percentage (%) (2) Parts per Million (ppm) & Parts per Billion (ppb) (3) Mole Fraction (X) (4) Molarity (M) (5) Molality (m)

  26. mass of solu A in solution total mass of solution mass of solu A in solution total mass of solution mass of solu A in solution total mass of solution Mass Percentage (%) (parts per Hundred)  100 or 102 Mass % of A = hundred 50. g of NaCl in 100.mL of H2O ppm = Parts per Million (ppm)  106 million 0.0005 g of NaCl in 100.mL of H2O Parts per Billion (ppb)  109 ppb = billion 0.000,000,5 g of NaCl in 100.mL of H2O

  27. moles of substance A total moles in solution (A+B+∙∙∙) XA = Mole Fraction (X) You can calculate the mole fraction of either the solvent or of the solute — make sure you find the one you need for your calculations! What are the mole fractions of methanol and water in a solution containing 128g CH3OH (MW=32.0) and 202 mL of H2O (MW=18.0)? Which is the solute and the solvent? Mole fraction of Solvent: Mole fraction of Solute:

  28. mol of solute L of solution M = Molarity (M) • Because volume is temperature dependent, molarity can change with temperature. {The volume of most substances expand with increasing temperature} What is the molarity of a solution containing 128g CH3OH (MW=32.0, d=0.791g/mL@25  C) and 202 mL of H2O (MW=18.0, d=1.00g/mL@25C)?

  29. moles of solute kg of solvent m = Molality (m) Because both moles and mass do not change with temperature, molality (unlike molarity) is not temperature dependent. What is the molality of a solution containing 128g CH3OH (MW=32.0, d=0.791g/mL@25  C) and 202 mL of H2O (MW=18.0, d=1.00g/mL@25C)?

  30. mol of solute L of solution mol of solute kg of solvent M = m = Changing Molarity to Molality If we know the density of the solution, we can calculate the molality from the molarity, and vice versa. Problem:What is the molality (m) of a 2.0M NaCl solution at 250C if its density is 1.2 g/mL? = 1.8 m NaCl

  31.  =MRT Colligative Properties • Properties that depend only on the number of solute particles present, not on the identity of the solute particles. • Vapor pressure lowering • 2. Boiling point elevation • 3. Freezing point • depression • 4. Osmotic pressure PA =XAPA Tb=Kb  m Tf=Kf m Pure solvent Solvent + Solute

  32. 1. Vapor Pressure (Pvap ) Lowering Because of solute-solvent intermolecular attraction, higher concentrations of nonvolatile solutesmake it harder for solvent to escape to the vapor phase. {PureSolv} Therefore, the vapor pressure of a solution (PA) is lower than that of the pure solvent (PA). PA =XAPA {*Solution}

  33. Raoult’s Law: {VapPress.WaterVsEthGlycol} PA = XA P A where • XA is the mole fraction of solvent A • P A is the normal vapor pressure of the pure solvent A at that temperature • P A is the new vapor pressure of solution at that temperature Consider: XA= 1 XA= 0.9 XA= 0.5

  34. 2. Freezing Point Depression & 3. Boiling Point Elevation f.pt. b.pt. f.pt. Solution f.pt. Pure solvent Nonvolatile solute-solvent interactions also cause solutions to have higher boiling points and lower freezing points than the pure solvent.

  35. Freezing Point Depression {f.pt. Equil} {f.pt. Lower} • The change in freezing point is proportional to the molality of the solution : Tf = Kf m • Here Kfis the molal freezing point depression constant of the solvent. Tf is subtracted from the normal freezing point of the solvent.

  36. 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.

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

  38. Diffusion Osmosis • The movement of particlesfrom an area of high concentration to lower concentration until a homogeneous solution has been formed. • The movement ofwater from an area of high concentration to lower concentration (diffusion) across asemipermeable membrane (SPM), until a homogeneous solution has been formed (equilibrium). The SPM allows smaller particles (water) to pass through, but blocks other larger particles. • Diffusion of water across a SPM.

  39. Osmosis In osmosis, there is net movement of solvent from the area of higher solvent concentration (lowersoluteconcentration) to the area of lower solvent concentration (highersoluteconcentration). Solvent Solution {*OsmoPress}

  40. {*OsmoSemiPerMemb} Solvent side (High Conc. H2O) Solution side (Low Conc. H2O) SPM

  41. Osmosis in Cells High Water H2O H2O • If the water concentration outside the cell is more than that inside the cell, the solution is hypotonic(low in solute). • Water will flow into the cell, and hemolysis results. (low solute) H2O H2O H2O H2O H2O H2O Low Water H2O H2O (high solute) H2O H2O H2O H2O H2O H2O H2O {EggOsmoPres}

  42. Osmosis in Blood Cells • If the water concentration outside the cell is lower than that inside the cell, the solution is hypertonic(high in solute). • Water will flow out of the cell, and crenation results. • Isotonic: equal concentrations. H2O (high solute) (salt water) H2O H2O H2O H2O H2O High Water H2O H2O (salt water) (low solute) H2O H2O H2O H2O H2O H2O (salt water) (salt water)

  43.  =()RT n V 3. Osmotic Pressure • The pressure required to stop osmosis, known as osmotic pressure, , is V = nRT  = MRT where M is the molarity of the solution

  44. Reverse Osmosis Reverse Osmosis {RevOsmo} {RevOsmo} Contaminated water Contaminated water Membrane Membrane H2O H2O H2O H2O H2O H2O H2O Pure

  45. Determining Molar Mass from Colligative Properties We can use the effects of a colligative property such as freezing point depression to determine the molar mass (MM) of a compound. A solution of an unknown nonvolatile electrolyte was prepared by dissolving 0.250g of the substance in 40.0 g of CCl4. The b. pt. of the resultant solution was 0.357 0C higher than that of the pure solvent. What is MM of solute? The Kb for CCl4 is 5.02 0C/m. Tb = Kb  m

  46. Determining Molar Mass from Colligative Properties We can use the effects of a colligative property such as osmotic pressure to determine the molar mass of a compound. The osmotic pressure of a protein solution was measured @ 250C to be 1.54 torr. The solution contained 3.50 mg of protein dissolved in water to make 5.00 mL of solution. Determine the molar mass (MM) of the protein.  = MRT

  47. Colligative Properties & Molar Mass • Vapor pressure lowering: • Boiling point elevation: • Freezing point depression: • Osmotic pressure: PA =XA PA = PA Tb = Kb  m = Kb Tf = Kf  m = Kf  =MRT =RT

  48. Colligative Properties of Electrolytes H2O C6H12O6 (s) C6H12O6(aq) Non-electrolyte: 1η 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. 1 η molecules H2O 3 η ions CaCl2 (s) Ca 2+(aq) + 2Cl-(aq) Electrolyte: 1η You would expect a 1 M solution of CaCl2 to show three times the change in freezing point that a 1 M solution of C6H12O6, however …………….

  49. van’t Hoff Factor One mole of NaCl in water does not really give rise to two moles of ions. 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. • Reassociation is more likely at higher concentration. • Therefore, the number of particles present is concentration dependent.

  50. =i MRT The van’t Hoff Factor • Equations • Vapor pressure lowering: • Boiling point elevation: • Freezing point depression: • Osmotic pressure: We modify the previous equations for Colligative Properties by multiplying them by the van’t Hoff factor, i PA = iXAPA Tb = iKb  m Tf =i Kf  m

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