Solutions and their Behavior

1. Solutions and their Behavior Goals: Calculate solution concentration. Describe the solution process. Apply colligative properties of solutions. Describe colloids and their applications.

2. Solutions A solution is a ______________ mixture of 2 or more substances in a single phase. One constituent is usually regarded as the SOLVENT and the others as SOLUTES.

3. Types of Solutions Solute Solvent Solution Example Gas Gas Gas Air (O2 in N2) ____ _____ _____ Club soda (CO2 in H2O) Liquid Liquid Liquid Wine (alcohol in H2O) ____ _____ _____ Saline sol. (NaCl in H2O) ____ _____ _____ 14-karat gold (Ag in Au) Aqueous solutions – those in which __________is the solvent.

4. Equilibrium Equilibrium – the ______ of the forward and reverse reactions are _______. Is __________ – reactants are changing to products, products to reactants; but at the same rate, there is no change in concentration of reactants or products.

5. Dynamic Equilibrium • Dynamic equilibrium – when the number of particles (ions or molecules) leaving the surface of the crystals is equal to the number that is returning. • The net quantity of particles in solution and that in undissolved crystals remain constant. • Saturated solution – a solution that contains all the solute that it can at equilibrium and at a given temperature. • 36 g NaCl per 100 g of H2O • Unsaturated solution – contains less than this quantity. • Any other with less: ex. 20 g NaCl per 100 g of H2O

6. Solutions • Solutions can be classified as saturated or _____________. • A saturated solution contains the ________ ______________ that dissolves at that temperature.

7. Dynamic Equilibrium • Precipitate – an insoluble or nearly insoluble solid that separates from a solution. • When a saturated solution is cooled, solute precipitates until the equilibrium is once again established at the lower temperature. 60oC 100 g water dissolve 95 g of lead (II) nitrate: Pb(NO3)2 5oC 100 g water dissolve 40 g of lead (II) nitrate The excess 55 g will separate as precipitate upon cooling, increasing the quantity of undissolved solute. • Supersaturated solution – a solution containing solute in excess of what it could contain if it were at __________. • It is not a stable system because it is not _______________. • Solute may precipitate when the solution is stirred or the inside of the container is scratched with a glass rod. • Addition of a “seed” crystal will nearly always result in __________ ___________________________________________________ __________________________________________________.

8. Solutions • Solutions can be classified as unsaturated or saturated. • A saturated solution contains the maximum quantity of solute that dissolves at that temperature. • ________________ SOLUTIONS contain more than is possible and are unstable.

9. Solutions • The solubility of a given solute depends on the relative __________________particles in the pure substances and in the solution. • Three things must happen: • The attractive forces holding the ions of the solute together must be overcome. • The attractive forces holding at least some of the solvent molecules must be overcome. • The solute and solvent molecules must interact; they must attract one another. • _____________ – process in which water molecules surround the solute ions.

10. Dissolving an Ionic Solid • The polarity of water molecules enables them to attract (and be attracted by) ions. Several ion-dipole interactions surround each ion and overcome the stronger ion-ion interactions.

11. Dissolving an Ionic Solid: Energetics • Almost all compounds of the Group 1A elements are soluble in water: NaCl, Na2SO4, K3PO4, LiBr • Many solids in which both ions are doubly or triply charged are essentially insoluble in water (forces holding the ions together are so strong that they cannot be overcome by the hydration of the ions): CaCO3, AlPO4, BaSO4.

12. Which ion is most strongly hydrated? • Na+ • Mg2+ • Cs+ Energy of hydration depends on the ________ of the ion and the _______ between the ion and the dipole.

13. Liquids dissolving Liquids • ___________– substances that can be mixed in all proportions (water and alcohol). • __________– the quantity that will dissolve is near zero (iron in water). • ___________– an appreciable quantity dissolves (sugar in water).

14. Solubility of Covalent Compounds • Like dissolves like. Nonpolar (or slightly polar) solutes dissolve best in nonpolar solvents; polar solutes dissolve best in polar solvents. • Water solubility of covalent compounds depends mainly on the ability of water to form _____________ to the solute molecules. Molecules containing a high proportion of nitrogen or oxygen atoms will dissolve in water.

15. Which of the following will dissolve in water? CH3OH Methyl alcohol CH3(CH2)2CH2OH Butyl alcohol CH3(CH2)10CH2OH Lauryl alcohol CH3CHO Acetaldehyde C12H22O11 Sucrose

16. Solubility of Gases Carbonated beverages (CO2 in water) Formalin (HCHO (formaldehyde gas)) Ammonia (NH3 in water) • Unlike most solids, gases become less soluble in water as the temperature _____________. The gas molecules acquire more kinetic energy and escape the solution. • At constant T, the solubility of a gas in water is directly proportional to the __________ of the gas in equilibrium with the aqueous solution. The higher the _______, the more gas will dissolve in a given volume of water. • The pressure inside soda bottles is high enough to dissolve the wanted CO2, once the bottle is opened, the pressure is released and the gas escapes.

17. Factors Affecting Solubility: Henry’s Law • The solubility of a gas in a liquid is directly proportional to the gas pressure. Sg = kHPg

18. Henry’s Law Gas solubility (mol/L) = kH • Pgas When Pgas drops, solubility drops. Think about: When would Henry’s Law not apply?

19. Factors Affecting Solubility: Le Chatelier’s Principle • A change in any of the factors determining an equilibrium causes the system to adjust so as to reduce or counteract the effect of the change. Henri Louis Le Chatelier (1884). – change in T, P, concentration of reactants or products. • Temperature affects solubility. • For all gases in water, solubility decreases with temperature.

20. Solubility and Temperature Simple correlations of solubility with structure or thermodynamic parameters are generally not successful.

21. Energetics of the Solution Process • If the enthalpy of formation of the solution is more negative that that of the solvent and solute, the enthalpy of solution is negative. • The solution process is exothermic!

22. Supersaturated Sodium Acetate • One application of a supersaturated solution is the sodium acetate “heat pack.” • Sodium acetate has an ENDOthermic heat of solution. Sodium acetate has an ENDOthermic heat of solution. NaCH3CO2 (s) ----> Na+(aq) + CH3CO2-(aq) Therefore, formation of solid sodium acetate from its ions is EXOTHERMIC. Na+(aq) + CH3CO2-(aq) ---> NaCH3CO2 (s) + heat + heat

23. Calculate Heat of solution (DHosoln) DHosoln = DHof products – DHof reactants

24. Colligative Properties On adding a solute to a solvent, the props. of the solvent are modified. • Vapor pressure decreases • Melting point decreases • Boiling point increases • Osmosis is possible (osmotic pressure) These changes are called COLLIGATIVE PROPERTIES. They depend only on the NUMBER of solute particles relative to solvent particles, not on the KIND of solute particles.

25. Concentration Units • An IDEAL SOLUTION is one where the properties depend only on the concentration of solute. • Need concentration units to tell us the number of solute particles per solvent particle. • The unit “molarity” does not do this!

26. Molarity and Molality • Molarity: Moles of solute per liter of solution. M = moles/L • Molality: Moles of solute per kilogram of solvent. m = moles/Kg Molality is Temperature _______________!

27. mol solute m of solute = kilograms solvent Concentration Units MOLALITY, m MOLE FRACTION, X For a mixture of A, B, and C WEIGHT % = grams solute per 100 g solution ppm = grams solute per 1 million g solution

28. What is the m of a 29.5% by mass ethanol solution (mw ethanol = 46g/mol)? Student should be familiar with concentration calculations.

29. Colligative Properties • Depend only on the _______________ ______________ relative to solvent particles, not on the KIND of solute particles • When a solute is present, the vapor pressure of the solvent is __________.

30. Liquid-Vapor Equilibrium • To understand colligative properties, study the LIQUID-VAPOR EQUILIBRIUM for a solution.

31. Raoult’s Law VP of H2O over a solution depends on the number of H2O molecules per solute molecule. Psolventproportional to Xsolvent Psolvent = Xsolvent • Posolvent Vapor Pressure of solvent over solution = (Mol frac solvent)•(VP pure solvent) RAOULT’S LAW

32. Raoult’s Law An _______ solution is one that obeys Raoult’s law. PA = XA • PoA Because mole fraction of solvent, XA, is always less than 1, then PA is always less than PoA. The vapor pressure of solvent over a solution is always LOWERED!

33. Changes in Boiling Points of Solvent See Figure 14.14

34. gas liquid Boiling Point Elevation A nonvolatile solute particle (purple) can block the escape of the solvent particles (blue) but has no effect on the return of the solvent particles from the vapor to the solution.

35. Boiling Point Elevation Elevation in BP = ∆TBP = KBP•m (where KBP is characteristic of solvent)

36. liquid solid Freezing Point Depression The rate at which solvent molecules (blue) leave the pure solid solvent is unaffected by the presence of solute particles (purple) nearby in the solution, but their rate of return to the solid is reduced.

37. Freezing Point Depression The freezing point of a solution is ________ than that of the pure solvent. FP depression = ∆TFP = KFP•m Ethylene glycol/water solution Pure water

38. Freezing Point Depression Water with and without antifreeze When a solution freezes, the solid phase is pure water. The solution becomes more concentrated.

39. Applications of DT • Antifreeze • Ice cream makers • CaCl2 on icy roads in winter • Measure molar masses • Distinguish between electrolytes and non-electrolytes.

40. Boiling Point Elevation and Freezing Point Depression ∆T = K•m•i A generally useful equation i = van’t Hoff factor = number of particles produced per formula unit. Compound Theoretical Value of i glycol 1 NaCl CaCl2

41. How much NaCl must be dissolved in 4.00 kg of water to lower the Freezing Point to -10.00 oC?

42. Which solution will have a lower freezing point? • 0.1 m glucose • 0.06 m NaCl • 0.06 m Na2SO4 Student should be familiar with predicting freezing and boiling point of solutions.

43. Osmosis Semipermeable membrane: solvent molecules can pass, but flow of solute is restricted. _________– net diffusion of water through a semipermeable membrane. Net flow of solvent from the more dilute solution (or pure solvent) into the more concentrated solution.

44. Osmotic Pressure Equilibrium is reached when pressure — the OSMOTIC PRESSURE, ∏ — produced by extra solution counterbalances pressure of solvent molecules moving thru the membrane. Solvent molecules move from pure solvent to solution in an attempt to make both have the same concentration of solute. Driving force is entropy

45. Osmotic Pressure • Osmotic Pressure (p) follows an equation much like the ideal gas law: p V = nRT or,p = MRT R = 0.0806 L atm/molK M = concentration of particles (moles or ions) in mol/L Useful for determining molar masses of large molecules like proteins and polymers.

46. Osmotic Pressure • Examples of osmosis are found in living organisms: Cells are semipermeable; their function and survival depend on maintenance of the same osmotic pressure inside the cell and outside in the extracellular fluid. • Isotonic solution – Iso-osmotic – a solution having the same osmotic pressure as body fluids (0.89% NaCl (mass/vol)).

47. Applications of Osmosis • ___________ solution – a solution having higher osmotic pressure than body fluids (larger than 0.89% NaCl). • ___________ solution – a solution having lower osmotic pressure than body fluids (less than 0.89% NaCl). Water flows out of the cell and cell wrinkles (crenation). Water flows into the cell and cell is swollen and may burst (plasmolysis). Normal cell

48. Applications of Osmosis

49. Applications of Osmosis • Reverse Osmosis Water desalination plant in Tampa

50. Dyalisis • Dialyzing membranes – membranes that pass small molecules and ions while holding back large molecules and colloidal particles. • Ions and small molecules always diffuse from higher concentration to lower concentration. • Dialysis – process in which small molecules and ions pass through a dialyzing membrane. In osmosis, osmotic membranes pass only solvent molecules. • Bags of cellophane or collodion. • Kidneys are a complex dialyzing system responsible for the removal of waste products from the blood. • Creatinine concentration > 900 mmol/L indication to start dialysis (kidney failure).