1. Properties of Solutions Chemical reactions often occur in _________ – a homogeneous mixture of two or more substances. The solute is the substance preset in a smaller amount, and the _______ is the substance present in a larger amount. In __________________, the solvent is always water. solutions solvent aqueous solutions

2. Properties of Solutions Solutions are not just a solid mixed in a liquid. That’s just one of many options. Solid in Liquid Liquid in _____ Gas in _____ Solid in ____ Gas in _____ Liquid in ____ Gas in _____ Liquid Solid Liquid Solid Gas Solid Gas

3. Spontaneous ! 2 nd law of Thermodynamics In order for ANYTHING to happen spontaneously ANYWHERE in the universe, the entropy of the universe must _________! Entropy is ________ or __________ --Which is more likely – your room to be neat or messy? **Messy is HIGH Entropy increase disorderrandomness

4. Aqueous Ionic Solution Properties _________– process in which an ion is surrounded by water molecules. For example: NaCl (s)  Na + (aq) + Cl - (aq) CaCl 2 (s)  _______________ H2OH2O H2OH2O Hydration Ca 2+ (aq) + 2 Cl - (aq)

5. Hydration

6. Ion-Dipole Forces ________________ are the electrostatic attraction between an ion and the oppositely charged end of a _____molecule. Example: Ions in water Interaction releases energy = (_______) of hydration =  H hyd Ion-dipole Forces polar enthalpy

7. Electrolytes Electrolyte - a substance that dissolves in water to form ions & make a solution that conducts _____________. Exp. Ionic like Na + Cl - or ______________ molecules Non electrolyte- _________ conduct electric current. Exp. Sugar and _________ molecules electric current highly polar does NOT non polar

8. Energetics of Solutions Intermolecular forces/interactions: 1) solute-solute 2) solvent-solvent 3) solute-solvent (_______ or _______) To form a solution, you must pull apart some of 1 and 2 (endo) but gain 3 (exo)!  H soln =  H 1 +  H 2 +  H 3  H soln overall can be endo or exo. solvationhydration

9. Energetics of Solutions Exothermic – feels ______ Solute + solvent  solution +_____ Endothermic – feels _____ Solute + solvent + heat  solution HOT heat cold

10. Energetics of Solutions If exo, new IMF will allow for the solution to be at __________ (energy is ______) solution will form! Feel: ___ If endo, it ___________ form solution. If too endo, the solute _____ dissolve. solute-solvent interactions not strong enough to break up 1, 2 interactions! Why do processes that are endo allow for solvation anyways? Can happen if the level of disorder ________ enough! Feel: ____ lower energy released hot may or may not won't increasescold

11. Liquid Solutes in Solvents “Like dissolves _____” is a useful rule to predict whether or not one substance dissolves in another So polar would dissolve with _____ Miscible: Substances that are soluble together in ____ proportions Immiscible: Substances that are ____ soluble with each other _____________ Like polar ALL NOT no matter what

12. Will it dissolve? “Like dissolves Like” is a useful rule to predict whether or not one substance dissolves in another How about water & oil? Water & vinegar? Water & grease? Gasoline & grease? Polar & nonpolar? Non polar & non polar? Immiscible Miscible Immiscible Miscible Immiscible Miscible

13. How does it look? Oil & Vinegar Water & Vinegar Non-polar & polar Immiscible Polar & Polar Miscible

14. Energetics of Solutions Solid solutes: When a solid is added to a liquid ______, molecules in solid phase overcome ____ and escape into solution, called __________ or __________. Opposite can also happen: molecules in solution collide with undissolved solid and come out of solution, called ______________. solvent IMF hydration.solvation. crystallization

15. Energetics of Solutions Solid solutes: Dynamic equilibrium is when rates of solute coming into and out of solution ___________. are equal

16. Factors affecting the rate of Dissolution - - Increasing __________________ --_________ a Solution -- _________ a solvent surface area of solute Agitating Heating

17. 3 Types of Solutions 1. Saturated solution Solution in equilibrium w/ undissolved solid. (there can be no increase in the # of dissolved solute particles) _______ measures the amount of _____ that is required to reach __________ for a given quantity of _______ (the max amount that can be dissolved!) e.g., NaCl in water: X g/100 ml H 2 O Solubilitysolute saturation solvent

18. 3 Types of Solutions 2. Unsaturated solution Has less solute than the ________ limit http://www.wwnorton.com/chemistry/overview/ch5.htm#precipitate 3. Supersaturated solution Under some conditions, you can get ____ solid dissolved than the solubility limit (_______, easy to get crystallization!) Examples: Honey & sodium thiosulfate solubility more unstable

19. Temp and Solubility For most solids, solubility  with __T For most gases, solubility  with __T Solutes that are Gases: Gases dissolve in liquids Solubility of a gas  as P ext __ Note: ______________ solubility largely independent of Pressure!    solid and liquid

20. Temp and Solubility

21. Temp and Gas Solubility As a gas heats up, due to _______ law its volume will __________ The volume is now ______ dense and will _______ & ________ liquid HEAT Escape!! Charles EXPAND LESS Riseescape The opposite will happen if you cool liquid – more can enter

22. Gas in atmosphere is in dynamic equilibrium w/ gas in solution: Rate at which gas particles strike solution surface is ______ at  P Pressure and Gas Solubility greaterso more enter

23. Pressure and Solubility Henry's Law: the solubility of a gas in a liquid is ________ proportional to the partial pressure of that gas on the surface of the liquid. _____________- the rapid escape of dissolved gas from its liquid solvent. So if you double pressure, you __________ solubility ! Think soda pop ! directly DOUBLE Effervescence

24. Solution Concentration Molarity – the number of moles of ______ in one liter of ________. Molarity (M) = amount of solute (mole) volume of solution (L) Ex. 0.50 mole NaOH = 0.50 M NaOH 1.00 L solutesolution

25. Solution Concentration Molality – the number of moles of solute in one kg of solvent. Molality (m) = amount of solute (mol) kg of solvent (kg) Ex. 0.12496 mol NaOH = 0.0625 m NaOH 2.00 kg H 2 O

26. Solution Concentration Mole fraction – the moles of component A divided by the total moles. Mole fraction(X) = moles component A total moles component Ex. 0.12496 mol NaOH 0.12496 mol NaOH + 110.98 mol H 2 O = 0.00112 mol fraction NaOH

27. Solution Concentration Weight % = mass % Weight % = mass component Ax 100 total mass solution 5.00 g NaOH = 0.249 % NaOH 5.00 g NaOH + 2000 g H 2 O Weight % usually reported in parts by mass (ppm, ppb, ppt, etc.)

28. Solution Concentration ppm = parts per million ppb = parts per billion A sample with 1.0 g of a substance in a sample with a total mass of 1.0 x 10 6 g would have a concentration of 1 ppm. one in a million! Ex. 5000 mg NaOH = 0.00248 = 2.49 ppm 2.01 L

29. Colligative Properties Colligative properties depend only on the number of solute molecules in a solution. Adding solute to a pure solvent cause: Decreased vapor pressure Boiling point (BP) elevation Freezing point (FP) depression Osmotic pressure

30. Vapor Pressure Vapor pressure of a solution is lower than that of pure solvent. Both liquids are in equilibrium w/ gas phase; particles escape more easily out of pure solvent. It has a higher vp!

31. Vapor Pressure Reduction in vp is determined mainly by the decrease in # of solvent particles at surface of a solution that can escape into the gas phase. This will be determined by mole fraction of the solvent. The amount by which vp is reduced in solution is proportional to the concentration of solute particles!

32. Vapor Pressure Consider a pure solvent and a solution:

33. Vapor Pressure In the pure solvent (mole fraction = 1) all surface particles are solvent particles. In the solution only some of the particles at the surface are solvent particles. High conc of solute or low mole fraction solvent  low vp Dilute conc of solute or high mole fraction solvent  vp is less reduced

34. Vapor Pressure

35. Boiling Point Elevation Boiling point (bp) occurs when vp = P ext Addition of of solute will  vp, and bp will occur at higher T for solution. why do we put salt in spaghetti water? let  T bp = increase in bp (it’s elevation!)  T bp = T bp (f, solution) – T bp (i, pure solvent) bp (solution) = bp (pure solvent) +  T bp

36. Boiling Point Elevation  T bp = k bp m wherek bp = molal bp elevation constant m = molality of solution depends on the solvent! Solvent norm FP ( o C) k fp ( o C/m) norm BP k bp ( o C/m) Water0.00 1.86100.000.52 Benzene5.5 5.12 80.12.53 Ethanol-117.3 1.99 78.41.22

37. BP Elevation Practice Pure H 2 O bp 100.0 o C and k bp = 0.52 o C/m What is bp elevation for a solution of 3.2 m CH 3 OH? What is the new bp? ->  T bp = (0.52 o C/m)(3.2 m) = 1.664 o C T bp final = T bp (pure) +  T bp = 100.0 o C + 1.664 o C = 101.7 o C

38. Freezing Point Depression Freezing point (fp) is also prop to molality. At the interface between liquid and solid (where solvent must solidify), there is less solvent, causing the system to release more heat to get the solvent particles to solidify. It has to get colder to actually freeze!

39.  T fp = k fp m k fp = molal fp depression constant  T fp is the decrease in fp  T fp = T fp (i, pure solvent) – T fp (f, solution) fp (soln) = fp (pure solvent) -  T fp Freezing Point Depression

40. FP Depression Practice for water k fp = 1.86 o C/m T fp (pure) = 0.0 o C What is the fp depression for a solution that is 3.2 m CH 3 OH? ->  T fp = (1.86 o C/m)(3.2 m) = 5.952 o C T fp (soln) = T fp (pure) –  T fp = 0.0 o C – 5.952 o C = -6.0 o C

41. Osmosis Osmosis is the net movement of solvent molecules from a less concentrated solution into a more concentrated one across a semi-permeable membrane.

42. Colligative Properties Colligative properties depend on # of particles. If solute is an electrolyte, must consider all of its aqueous ions! When performing calculations, multiply the concentration by the van't Hoff factor (i) i = the # of ions after dissociation/formula unit. NaOH (aq)  Na + (aq) + OH - (aq) i = 2 MgCl 2 (aq)  Mg 2+ (aq) + 2Cl - (aq) i = 3

43. Colligative Properties  T bp = ik bp m  T fp = ik fp m  = iMRT The result is that, for example: 1.0 M CH 3 OH (aq) = 1.0 M particles 0.5 M NaOH (aq)  0.5 M Na + (aq) + 0.5 M OH - (aq) = 1.0 M particles Both solutions have same reduction in vapor pressure!

44. BP Elevation Practice 401.2 g solute were dissolved in 2.0 kg of water. The bp elevation of the solution is 2.3 o C. What is the molar mass of the solute?hint:  T bp = k bp m -> 2.3 o C = (0.52 o C/m)(m) -> m = 4.423 m 4.423 m = (4.423 mol /kg solvent)(2.0 kg) = 8.846 mol solute molar mass: 401.2 g /8.846 mol solute = 45.35 g/mol

45. Solution types Homogenous Solution – looks the same through-out, cannot separate by filtering Heterogenous Solution – CAN separate out by filtering. May settle out. Suspensions: A heterogenous mixture in which droplets or particles are suspended in a liquid

46. Colliods Solutions containing suspensions of particles that will not "settle out" (intermediate between solution and suspension) particles are large, have large surface area.

47. Colliods Colloids in water can be … Hydrophobic = water "fearing;“ usually nonpolar or polar w/ weak attractions to water. Hydrophilic = water "loving;“ usually polar(e.g., proteins)

48. Colliods Emulsions are colloidal dispersions of one liquid in another that would normally be immiscible (e.g., mayonnaise, hand lotion) Usually contain an emulsifying agent (soap or a protein – bridges between the different types of molecules)

49. Colliods Surfactant/detergent are used as soaps that have polar and nonpolar ends – can "bridge" between polar aqueous solution and nonpolar grease/dirt.

50. Energetics of Solutions Liquids solutes: Liquids are called miscible if they mix in any proportions (water and ammonia) and immiscible if they will not mix in any proportions (water and oil). “Like dissolves like“ Polar solutes dissolve in polar solvents. Nonpolar solutes dissolve in nonpolar solvents. (grease in gasoline)