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Chapter 14

Chapter 14. Solutions. Overview. Solution formation Types of solution Solubility and the solution process Effect of temperature and pressure on solubility Ways of expressing concentration Colligative properties Vapor pressure of a solution

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Chapter 14

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  1. Chapter 14 Solutions

  2. Overview • Solution formation • Types of solution • Solubility and the solution process • Effect of temperature and pressure on solubility • Ways of expressing concentration • Colligative properties • Vapor pressure of a solution • Boiling-point elevation and freezing point depression. • Osmosis • Colligative properties of ionic solutions

  3. Types of Solution • Solution – homogeneous mixture of two or more substances of ions or molecules. E.g. NaCl (aq) • Colloid – appears to be a homogeneous mixture, but particles are much bigger, but not filterable. E.g. Fog, smoke, whipped cream, mayonnaise, etc. • Suspension: larger particle sizes, filterable. E.g. mud, freshly squeezed orange juice.

  4. Solution Composition Solvent = component which is the component in greater amount. Solute = component which is present in the smaller amount. Aqueous Solutions = aqueous solutions are those in which water is the solvent.

  5. Solution Composition • Gaseous = gases are completely miscible in each other. • Liquid = gas, liquid or solid solute dissolved in solvent. • Solid = mixture of two solids that are miscible in each other to form a single phase.

  6. Electrolytes • Electrolytes are solutes that dissolve in water to form ions and consequently are capable of conducting electricity. Electrolytes consist of ionic compounds as well as binary molecules of hydrogen (acids). • Examples: • NaCl(s) ----------> Na+(aq) + Cl-(aq) • HF(g) ------------> H+(aq) + F-(aq) • Ca(NO3)2(s) ----------> Ca2+(aq) + 2NO3-(aq)

  7. Nonelectrolytes • Nonelectrolytes are solutes that dissolve in water to form molecules and consequently are incapable of conducting electricity. • CH3OH(l) ----------> CH3OH(aq)

  8. Concentration of Solutes - Qualitative • Dilute - this means that the solution contains only a small amount of solute • Concentrated - this means that the solution contains more solute

  9. As you may have guessed - these are not real good ways of describing solution concentrations. But they do occasionally serve useful purposes. For example, in the case of acids and bases the terms dilute and concentrated refer to special concentrations of the solutions: • Solution Dilute Concentrated • HCl 6.00 M 12.0 M • HNO3 6.00 M 15.0 M • H2SO4 3.00 M 18.0 M • NH3 6.00 M 15.0 M

  10. Solubility and the Solution Process • The solid dissolves rapidly at first but as the solution approaches saturation the net rate of dissolution decreases since the process is in dynamic equilibrium. • When the solution has reached equilibrium the amount of solute does not change with time; • At equilibrium: the rate of dissolution = rate of solution Fig. 12.2 Solubility Equilibrium

  11. http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/molvie1.swfhttp://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/molvie1.swf

  12. Concentration of Solutes Using Equilibria • Concentrations in terms of equilibria rely heavily on the solubility of the solute in the particular solvent. Solubility is defined as the maximum amount of solute that will dissolve in a particular solvent under the specified conditions of temperature and pressure. Solubility is generally given in terms of grams of solute per 100 grams of solvent.

  13. Saturated solution: a solution in which a dynamic equilibria exists between the undissolved solute and the solution. The solution contains the maximum amount of dissolved solute according to solubility.

  14. Unsaturated solution: a solution that contains less than the maximum amount of dissolved solute.

  15. Supersaturated: a solution that contains a greater concentration of solute than a saturated solution. This type of solution is unstable. Generally, if you add a tiny crystal of the solute to the solution all the excess solute will very rapidly come out of solution.

  16. Molarity, M Molarity is defined as the moles of solute per liter of solution. (Chapter 5 - see pages 228 - 235 if you can't remember.)

  17. Molality, m Molality is defined as the moles of solute per kilogram of solvent. This is new to this chapter and will be very important when determining certain properties of solutions.

  18. Determine the molality of a solution prepared by dissolving 1.44 g NaCl into exactly 100.0 mL of water. Assume the density of water is 1.00 g/mL.

  19. Mole Fraction, X You can have a mole fraction of the solute or a mole fraction of the solvent. Mole fraction is defined as the moles of solute divided by the total number of moles or the moles of solvent divided by the total number of moles. The mole fraction is expressed as a decimal. An explanation of this can also be found on page 656.

  20. Determine the mole fraction of a solution made from dissolving 30.0 g H2O2 with 70.0 g H2O.

  21. Weight percent/PPM/PPB Weight percent is defined as the mass of component of the solution divided by the total mass of the solution - multiplied by 100. Parts per million/parts per billion are determined by dividing the mass of the solute by the total mass and then multiplying of 1 million or 1 billion.

  22. Determine the mass % of a solution prepared by dissolving 1.44 g of NaCl in 100.0 mL of water. Assume that the density of water is 1.00 g/mL

  23. Determine mass % of solution made from dissolving 30.0 g H2O2 with 70.0 g H2O.

  24. Determine molality of 30% H2O2(aq)

  25. Concentrated ammonia is 14.8 M and has a density of 0.900 g/mL. What is the molality?

  26. Factors Affecting Solubility • Nature of the Solute and Solvent • Temperature • Pressure

  27. Mixing of Gas Molecules Nature of the Solute and Solvent • Gases = generally completely soluble in each other because of entropy (tendency towards maximum randomness). • Molecules in gas phase are far apart from each other and not interacting strongly with each other in solution.

  28. Electrolytes and Nonelectrolytes • Most nonelectrolytes that are appreciably soluble in water are hydrogen bonded (CH3OH, H2O2, sugars) • Other types of nonelectrolytes are generally more soluble in nonpolar or slightly polar solvents (C6H6, CCl4)

  29. Molecular Solutions • Molecular compounds with similar chemical structures and polarities tend to be miscible. • Homologous alcohol series have polar and non-polar ends.

  30. Ionic Solutions • Solubility affected by: • Energy of attraction (due Ion-dipole force) affects the solubility. Also called hydration energy, • Lattice energy (energy holding the ions together in the lattice. Related • to the charge on ions; larger charge means higher lattice energy. • Inversely proportional to the size of the ion; large ions mean smaller lattice energy.

  31. Solubility increases with increasing ion size, due to decreasing lattice energy; Mg(OH)2(least soluble), Ca(OH)2, Sr(OH)2, Ba(OH)2(most soluble) (lattice energy changes dominant). • Energy of hydration increases with for smaller ions than bigger ones; thus ion size. MgSO4(most soluble),... BaSO4 (least soluble.) Hydration energy dominant.

  32. Heat of Solution and Solubility • Hsoln is sometimes negative and sometimes positive. • Solvent – solvent interactions: energy required to break weak bonds between solvent molecules. • Solute – solute interactions: energy required to break intermolecular bonds between the solute molecules. • Solute – solvent interactions: H is negative since bonds are formed between them.

  33. Energy Changes and Solution Formation

  34. Effect of Temperature on Solubility • An increase in temperature favors the endothermic process • Solid + water -----------> Solution; DH for these solutions is usually endothermic so generally an increase in temperature will increase solubility. 3. Gas + water ------------> Solution; DH for these solutions is usually exothermic (WHY?) so generally an increase in temperature will decrease solubility.

  35. Solubility: Temperature Dependence • All solubilities are temperature dependent; must report temperatures with solubilities. • Temperature related to sign of Hsoln;

  36. Predict the temperature dependence of the solubility of Li2SO4, Na2SO4 and K2SO4 if their Hsoln are 29.8 kJ/mol, 2.4 kJ/mol and +23.8 kJ/mol, respectively.

  37. Solubility: Pressure Dependence • Pressure has little effect on the solubility of a liquid or solid, but has dramatic effect on gas solubility in a liquid. • Henry’s law S = kHP. Allows us to predict the solubility of a gas at any pressure.

  38. At 25C P(O2 in air) = 0.21 atm. Its solubility in water is 3.2x10-4M. Determine its solubility when pressure of O2 = 1.00 atm.

  39. Physical Behavior of Solutions: Colligative Properties • Compared with the pure solvent the solution’s: • Vapor pressure is lower • Boiling point is elevated • Freezing point is lower • Osmosis occurs from solvent to solution when separated by a membrane.

  40. Vapor-Pressure Lowering of Solutions: Raoult’s Law • Raoult’s Law: Psoln = PsolvxXsolv • Non–volatile solute: vapor pressure decreases upon addition of solute. • Linear for dilute solutions • http://dwb4.unl.edu/ChemAnime/solutions.htm

  41. Vapor pressure lowering : P = Po P = Po(1Xsolv) Determine vapor pressure lowering when 5.00 g of sucrose added to 100.0 g of H2O. MM(sucrose) = 342.3 g/mol. The vapor pressure of water at 25°C is 23.8 mmHg.

  42. Determine the mass of sucrose dissolved in 100.0 g of water if the vapor pressure was 20.0 mmHg.

  43. BP Elevation of Solutions The magnitude of the change in FP and BP is directly proportional to the concentration of the solute (molality) – expressed in terms of the total number of particles in the solution. • BP Elevation The magnitude of the BP increase is given by the equation: where Kb has units of °Ckg/mol or °C/m

  44. Determine boiling point elevation when 5.00 g of sucrose is added to 100.0 g of H2O. MM(sucrose) = 342.3 g/mol. Kb = 0.521 C/m.

  45. FP Depression of Solutions • FP Depression: linear variation with composition and given by: where the units for this constant are the same as for Kb Determine freezing point depression when 5.00 g of sucrose is added to 100.0 g of H2O. FM(sucrose) = 342.3 g/mol. Kf = 1.86°C/m.

  46. Osmosis and Osmotic Pressure Osmosis: the passage of solvent through a membrane from the less concentrated side to the more concentrated side. Osmotic pressure: the amount of pressure necessary to stop Osmosis. • Small molecules such as water can move through certain types of materials (membranes). • The tendency for this to occur is related to the molarity of the solution, is also a function of the temperature and is measured with a device called a Thistle tube. where M = is molarity of solute particles

  47. Determine osmotic pressure of a solution containing 0.100 g of hemoglobin (molecular mass = 6.41x104 amu) in 0.0100 L at 1.00C.

  48. Reverse Osmosis • Application of a pressure to the solution (that is equal to or greater than the Osmotic pressure) and the solvent flows from the more concentrated side to the other one. • This process is used to obtain pure water from salt water.

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