I. Types of Mixtures

1. Ch. 14 – Mixtures & Solutions I. Types of Mixtures

2. A. Definitions • Mixture = Variable combination of 2 or more pure substances • Homogeneous = uniform composition throughout • Heterogeneous = variable composition Heterogeneous Homogeneous

3. A. Definitions • Solution – homogeneous mixture Solute – substance being dissolved Solvent – dissolves the solute

4. B. Mixtures • Gases can also mix with liquids • Gases are usually dissolved in water • Examples are carbonated drinks • Homogeneous mixtures (solutions) • Contain sugar, flavorings and carbon dioxide dissolved in water

5. Tyndall Effect B. Mixtures • Solution • homogeneous • very small particles • no Tyndall effect • particles don’t settle • Ex: rubbing alcohol

6. B. Mixtures • Colloid • heterogeneous • medium-sized particles • Tyndall effect :the scattering of light by colloidal particles. • particles don’t settle • Ex: milk

7. B. Mixtures • Suspension • Heterogeneous • large particles usually > 1000nm • particles settle • Tyndall effect • Ex: fresh-squeezed lemonade

8. B. Mixtures • Examples: • mayonnaise • muddy water • fog • saltwater • Italian salad dressing • colloid • suspension • colloid • true solution • suspension

9. Ch. 14 –Solutions II. Factors Affecting Solvation(p. 489 – 496)

10. A. Solvation • Solvation – the process of dissolving solute particles are surrounded by solvent particles First... solute particles are separated and pulled into solution Then...

11. Solution Process • For a solute to be dissolved in a solvent, the attractive forces between the solute and solvent particles must be great enough to overcome the attractive forces within the pure solvent & pure solute. • The solute & the solvent molecules in a solution are expanded compared to their position within the pure substances.

12. B. Solvation • Dissociation • Separation/Solvation of an ionic solid into aqueous ions • For ionic solids, the lattice energy describes the attractive forces between the solute molecules (i.e. ions) • For an ionic solid to dissolve in water, the water-solute attractive forces has to be strong enough to overcome the lattice energy NaCl(s)  Na+(aq) + Cl–(aq)

13. B. Factors Affecting Solvation • Molecules are constantly in motion according to… • Kinetic Theory • When particles collide, energy is transferred

14. B. Factors Affecting Solvation • Solubility = max. amount of a solute that will dissolve in a solvent @ a specific T • Smaller pieces of a substance dissolve faster b/c of larger surface area • Stirring or shaking speeds dissolving b/c particles are moving faster and colliding more • Heating speeds dissolving of solids • Not all substances dissolve

15. C. Solubility • Water is universal solvent b/c of its polarity • If something can dissolve in something else, it is said to be soluble or miscible • If it cannot dissolve, it is said to be insoluble or immiscible • “Like dissolves like”

16. NONPOLAR NONPOLAR POLAR POLAR C. Solubility “Like Dissolves Like”

17. Saturated Solutions • A solution that can contain the maximum amount of solute at a given temperature (if the pressure is constant). • Solution said to be at a dynamic equilibrium • Any point on the line • Ex: At 90 ° C 40 g of NaCl (s) in 100 g of H2O represent a saturated solution

18. Unsaturated Solution • A solution that can contain less than the maximum amount of solute at a given temperature (if the pressure is constant). • It is any value under the solid line on the solubility graph

19. Supersaturated Solutions • A solution that can contain greater than the maximum amount of solute at a given temperature (if the pressure is constant). • A supersaturated solution is very unstable & the amount of solute in excess can precipitate or crystallize. • It is any value above the solid line on the solubility graph

20. UNSATURATED SOLUTION more solute dissolves SATURATED SOLUTION no more solute dissolves SUPERSATURATED SOLUTION becomes unstable, crystals form Solubility concentration

21. Solubility Any solution can be made Saturated, Unsaturated, or Supersaturated by changing the Temperature.

22. C. Solubility • Solubility Curves • maximum grams of solute that will dissolve in 100 g of solvent at a given temperature • varies with temp • based on a saturated soln

23. C. Solubility • Solubility Curve • shows the dependence of solubility on temperature

24. C. Solubility • Solids are more soluble at... • high temperatures. • Gases are more soluble at... • low temperatures • high pressures (Henry’s Law). • With larger mass (LDF) • EX: nitrogen narcosis, the “bends,” soda

25. - + - - + + acetic acid salt sugar • Although H2O is a poor conductor of electricity, dissolved ions in an aqueous solution can conduct electricity. • Ionic aqueous solutions are known as electrolytes. Non- Electrolyte Weak Electrolyte Strong Electrolyte solute exists as molecules only solute exists as ions and molecules solute exists as ions only DISSOCIATION IONIZATION

26. Ch. 16 – Solutions II. Solution Concentration(p. 480 – 486)

27. A. Concentration • The amount of solute in a solution • Describing Concentration • % by mass - medicated creams • % by volume - rubbing alcohol • ppm, ppb - water contaminants • molarity - used by chemists • molality - used by chemists

28. substance being dissolved total combined volume B. Percent Solutions • Percent By Volume (%(v/v)) • Concentration of a solution when both solute and solvent are liquids often expressed as percent by volume

29. B. Percent Solutions • Find the percent by volume of ethanol (C2H6O) in a 250 mL solution containing 85 mL ethanol. • Solute = 85 mL ethanol • Solution = 250 mL 85 mL ethanol 250 mL solution % (v/v) = = 34% ethanol (v/v) x 100

30. substance being dissolved total combined volume C. Molarity • Concentration of a solution most often used by chemists

31. C. Molarity 2M HCl What does this mean?

32. D. Molarity Calculations LITERS OF GAS AT STP Molar Volume (22.4 L/mol) MASS IN GRAMS NUMBER OF PARTICLES MOLES Molar Mass (g/mol) 6.02  1023 particles/mol Molarity (mol/L) LITERS OF SOLUTION

33. D. Molarity Calculations • How many moles of NaCl are required to make 0.500L of 0.25M NaCl? 0.500 L sol’n 0.25 mol NaCl 1 L sol’n = 0.013 mol NaCl

34. D. Molarity Calculations • How many grams of NaCl are required to make 0.500L of 0.25M NaCl? 58.44 g NaCl 1 mol NaCl 0.500 L sol’n 0.25 mol NaCl 1 L sol’n = 7.3 g NaCl

35. D. Molarity Calculations • Find the molarity of a 250 mL solution containing 10.0 g of NaF. 10.0 g NaF = 0.24 mol NaF 1 mol NaF 41.99 g NaF 0.24 mol NaF = 0.95 M NaF 0.25 L

36. E. Dilution • Preparation of a desired solution by adding water to a concentrate • Moles of solute remain the same

37. E. Dilution • What volume of 15.8M HNO3 is required to make 250 mL of a 6.0M solution? GIVEN: M1 = 15.8M V1 = ? M2 = 6.0M V2 = 250 mL WORK: M1 V1 = M2 V2 (15.8M)V1 = (6.0M)(250mL) V1 = 95 mL of 15.8M HNO3

38. mass of solvent only 1 kg water = 1 L water F. Molality

39. G. Molality Calculations • Find the molality of a solution containing 75 g of MgCl2 in 250 mL of water. 75 g MgCl2 1 mol MgCl2 95.21 g MgCl2 = .79 mol MgCl2 .79 mol MgCl2 = 3.2m MgCl2 0.25 kg water

40. G. Molality Calculations • How many grams of NaCl are req’d to make a 1.54m solution using 0.500 kg of water? 0.500 kg water 1.54 mol NaCl 1 kg water 58.44 g NaCl 1 mol NaCl = 45.0 g NaCl

41. 500 mL of 1.54M NaCl 500 mLwater 500 mL volumetric flask 500 mL mark 45.0 gNaCl H. Preparing Solutions • 1.54m NaCl in 0.500 kg of water • mass 45.0 g of NaCl • add water until total volume is 500 mL • mass 45.0 g of NaCl • add 0.500 kg of water

42. 95 mL of15.8M HNO3 250 mL mark water for safety H. Preparing Solutions • 250 mL of 6.0M HNO3by dilution • measure 95 mL of 15.8M HNO3 • combine with water until total volume is 250 mL • Safety: “Do as you oughtta, add the acid to the watta!” or AA – add acid!

43. Ch. 14 – Mixtures & Solutions IV. Colligative Properties of Solutions(p. 498 – 504)

44. A. Definition • Colligative Property • property that depends on the concentration of solute particles, not their identity • Examples: vapor pressure, freezing point, boiling point

45. B. Types

46. B. Types • Freezing Point Depression (Tf) • f.p. of a solution is lower than f.p. of the pure solvent • Boiling Point Elevation (Tb) • b.p. of a solution is higher than b.p. of the pure solvent • Vapor Pressure Lowering • lower number of solvent particles at the surface of the solution; therefore, this lowers the tendency for the solvent particles to escape into the vapor phase.

47. B. Types • Applications • salting icy roads • making ice cream • antifreeze • cars (-64°C to 136°C)

48. C. Calculations T:change in temperature (°C) i: Van’t Hoff Factor (VHF), the number of particles into which the solute dissociates m: molality (m) K:constant based on the solvent (°C·kg/mol) or (°C/m) T = i · m · K

49. C. Calculations •  T • Change in temperature • Not actual freezing point or boiling point • Change from FP or BP of pure solvent • Freezing Point (FP) •  TF is always subtracted from FP of pure solvent • Boiling Point (BP) •  TB is always added to BP of pure solvent

50. C. Calculations • i – VHF • Nonelectrolytes (covalent) • remain intact when dissolved • 1 particle • Electrolytes (ionic) • dissociate into ions when dissolved • 2 or more particles