Solutions

1. Solutions

2. Solutions • A solution is formed when one substance disperses uniformly throughout another. • Intermolecular forces function between solute particles and the solvent molecules that surround them.

3. Solutions form when the attractive forces between solute and solvent are comparable in magnitude with those that exist between the solute particles themselves or between the solvent particles themselves. • (see pgh 2 on pg 470 B&L)

4. When separated ions are surrounded by water molecules it is called solvation. • When the solvent is water, the interactions are known as hydration.

5. Occur in all phases • The solvent does the dissolving. • The solute is dissolved. • There are examples of all types of solutes dissolving in all types of solvents. • We will focus on aqueous solutions.

6. Ways of Measuring Concentration • Molarity • Molality • % mass • Normality (just read this section) • Mole Fraction

7. Molarity • Molarity = moles of solute Liters of solvent • Units are moles liter • Use the symbol “M” for molarity

8. Molarity • Try some molarity calculations

9. Molality • Molality = moles of solute Kilograms of solvent • Units are mole/Kg • Molality is abbreviated “ m “

10. Molality • Try some molality calculations

11. % Mass • % mass = Mass of solute x 100 Mass of solution

12. Normality • Normality - read but don’t focus on it.

13. Mole Fraction • Mole fraction of component A cA = NANA + NB

14. Dilution • M1 V1 = M2 V2 • “M” is molarity • “V” is volume (in liters)

15. The overall enthalpy change in forming a solution ∆Hsoln, is the sum of the following: • ∆Hsoln = ∆H1 + ∆H2 + ∆H3

16. Energy of Making Solutions • Heat of solution ( DHsoln ) is the energy change for making a solution. • Most easily understood if broken into steps. • 1.Break apart solvent • 2.Break apart solute • 3. Mixing solvent and solute

17. 1. Break Apart Solvent • Have to overcome attractive forces, this requires energy. DH1 >0 (endo) 2. Break apart Solute. • Have to overcome attractive forces. DH2 >0 (endo) the solvent molecules need to separate to make room for the solute

18. 3. Mixing solvent and solute • DH3 depends on what you are mixing. • Molecules can attract each other DH3 is large and negative (∆H < 0 , exo) . • Molecules can’t attract DH3 issmall and negative. • This explains the rule “Like dissolves Like”

19. Solute and Solvent DH3 DH2 Solvent DH1 DH3 Solution Solution • Size of DH3 determines whether a solution will form Energy Reactants

20. Types of Solvents and solutes • If DHsoln is small and positive, a solution will still form because of entropy. • There are many more ways for them to become mixed than there is for them to stay separate.

21. The formation of a solution can be either exo or endothermic. • If the process is exo it will tend to proceed spontaneously. • A solution will not form if ∆Hsoln is too endothermic.

22. The solute/solvent interactions must be strong enough to make ∆H3 comparable in quantity with ∆H1 + ∆H2 • Read pgh 2 and 3 on pg 472 B&L • Two factors are involved in processes that occur spontaneously: * energy * disorder

23. Processes in which the energy content of the system decreases tend to occur spontaneously. (spontaneous processes tend to be exothermic)

24. Types of Solutions • Solid solution- Alloys are the most common solid solutions containing 2 or more metals • Liquid solution- Miscible - 2 or more liquids that can mix in any amount Immiscible - liquids that cannot mix in any proportions • Aqueous solutions have water as the solvent

25. Saturated Solutions and Solubility • Several types of solutions: • Unsaturated • Saturated • Supersaturated

26. Concentrations of Solutions • Unsaturated -- A solution that contains less than the maximum amount of solute that can be dissolved at that temperature. • Saturated Solution -- A solution containing the maximum amount of solute that can be dissolved at that temperature.

27. Supersaturated Solutions? • Supersaturated -- A solution that contains more solute than would normally dissolve at that temp. Unstable! • How can a solution be supersaturated? • Well, how can we dissolve MORE solute? • Heat! • So, heat a solution, dissolve MORE solute, then cool it CAREFULLY.

28. A formerly supersaturated solution -- a single crystal of the solute introduced will cause ALL of the excess solute to come out of solution suddenly! http://www.chem.ufl.edu/~itl/2045/lectures/lec_i.html

29. Factors Affecting Solubility • Temperature • Pressure • Surface Area • Agitation

30. Factors Affecting Solubility Pressure • External pressure has no effect on the solubility of solids or liquids because solids and liquids are not appreciably compressed when pressure is increased. • Gases can be compressed easily.

31. Pressure - Solubility • Gas compression increases the frequency with which gas molecules hit the liquid phase and enter it, thereby increasing the solubility. • The effect of pressure on the solubility of gases is expressed in Henry’s law: solubilitygas = k Pgas

32. Henry’s law solubilitygas = k Pgas • Solubility is expressed as molarity. • “k” is the Henry’s law constant • Pgas is the partial pressure of the gas.

33. Pressure effects • Changing the pressure doesn’t effect the amount of solid or liquid that dissolves • They are incompressible. • Pressure does effect gases.

34. Dissolving Gases • Pressure effects the amount of gas that can dissolve in a liquid. • The dissolved gas is at equilibrium with the gas above the liquid.

35. The gas is at equilibrium with the dissolved gas in this solution. • The equilibrium is dynamic.

36. If you increase the pressure the gas molecules dissolve faster. • The equilibrium is disturbed.

37. The system reaches a new equilibrium with more gas dissolved. • Henry’s Law. P= kC Pressure = constant x Concentration of gas

38. Factors Affecting Solubility • Temperature • As a general rule, increasing temperature drives molecules toward the more random phase. Therefore, an increase in temperature usually increases the solubility of solids in a liquid and always decreases the solubility of gases in a liquid. • Decreasing temperatures have the opposite effect.

39. Temperature Effects • Increased temperature increases the rate at which a solid dissolves. • We can’t predict whether it will increase the amount of solid that dissolves. • We must read it from a graph of experimental data.

40. The next slide shows Solubility Curves for several substances.

41. 100 40 60 80 20

42. Gases are predictable • As temperature increases, solubility decreases. • Gas molecules can move fast enough to escape.

43. Colligative Properties • Colligative properties depend only on the number of dissolved particles in solution and not on their identity. • The four Colligative Properties are: • Vapor pressure reduction • Boiling point elevation • Freezing point depression • Osmotic pressure

44. Vapor Pressure • A liquid in a closed container will establish an equilibrium with its vapor. • When that equilibrium is reached, the pressure exerted by the vapor is called the vapor pressure.

45. A substance that has no measurable vapor pressure is called non-volatile. • A substance that has a vapor pressure is called volatile.

46. Vapor Pressure of Solutions • A non-volatile solvent lowers the vapor pressure of the solution. • The molecules of the solventmust overcome the force of both the other solvent molecules and the solute molecules.

47. Vapor Pressure • http://www.unit5.org/christjs/Vapor_Pressure_Boiling_Point.htm A                                B In container A, the liquid is evaporating.  Some of the molecules have enough kinetic energy to escape (turn to a gas) by pushing against the pressure of the atmosphere.  Container B shows the flask is saturated.  When new molecules of liquid are vaporized, the gas cannot hold additional molecules, therefore some of the molecules condense back to liquid.

48. Raoult’s Law: • Psoln = csolvent x Psolvent • Psoln is the vapor pressure of the solution • csolvent is the mole fraction of the solvent • Psolvent is the vapor pressure of the pure solvent • This applies only to an ideal solution where the solute doesn’t contribute to the vapor pressure.

49. Raoult’s Law • An ideal gas obeys the ideal gas law, and an ideal solution obeys Raoult’s Law. • Real solutions best approximate ideal behavior when the solute concentration is low and when the solute and solvent have similar molecular sizes and similar types of intermolecular attractions.

50. Water has a higher vapor pressure than a solution Aqueous Solution Pure water