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Ch. 14/15 – Solids, Liquids and Solutions

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1 Ch. 14/15 – Solids, Liquids and Solutions

2 Intermolecular Forces
Forces of attraction between different molecules rather than bonding forces within the same molecule. Dipole-dipole attraction Between particles with charged sides ml Hydrogen bonds - Strongest Between H’s & F, O, N on other molecules Dispersion forces - Weakest Caused by electrons shifting towards one end of a molecule.

3 “Water: The Magnificent Dipole”
One side of water is negatively charged because the oxygen atom keeps the shared electrons longer than the hydrogen atoms. As a result the oxygen side is negatively charged and the hydrogen side of water is positively charged. O

4 It is its polarity and hydrogen bonding that give water its many unusual Properties! i.e. high boiling point, expansion upon freezing, and surface tension

5 Hydrogen bonds - Strongest Between H’s & F, O, N on other molecules
Ice Liquid

6 Water is always trying to pull itself into a tight ball as long as there is nothing nearby that has a charge on it. Therefore, this surface is not repelling water; it’s simply not attracting it and keeping water from doing what it does naturally. Water pulls on itself so much that it forms a “skin.” It’s called surface tension.

7 Wax does not repel water
We’ve heard that wax or oils repel water. But that isn’t true. Water is so attracted to other water molecules that anything between them is squeezed out of the way. O O Oil droplet O O O

8 Forces and Phases Substances with very little intermolecular attraction exist as gases Substances with strong intermolecular attraction exist as liquids Substances with very strong intermolecular (or ionic) attraction exist as solids

9 “V” is for Vocabulary! Vapor – gaseous state of a substance that is not normally a gas at room temperature. Volatile – evaporates rapidly (due to weak intermolecular forces) Ex. ‘thin’ liquids Viscous – evaporates slowly (due to strong intermolecular forces) Ex. ‘thick’ liquids

10 Phase Differences

11 Three Phases of Matter

12 Liquids A decrease in the average kinetic energy of gas particles causes the temperature to decrease. As it cools, the particles tend to move more slowly, if they slow down enough, attractive forces - called van der Waal’s forces –pull them very close together so they can only slip & slide past each other. Condensation – Change of a gas to a liquid It is now in liquid form!

13 The Nature of Liquids The conversion of a liquid to a gas or vapor below its boiling point is called Vaporization (occurs at the surface of a liquid) In an open container, this process is called Evaporation Particles near the surface with enough kinetic energy that happen to bounce in the right direction escape!

14 Microscopic view of a liquid near its surface.

15 The Nature of Liquids Eventually the particles will lose energy and return to the liquid state, or condense. What are the odds that they will return to the original liquid? What if we cover the container? So, the particles begin to evaporate, then some begin to condense. Eventually, the number of particles evaporating will equal the number condensing & the space above the liquid will be saturated with vapor

16 A dynamic equilibrium now exists where
Rate of evaporation = rate of condensation Note that there will still be particles that evaporate and condense But, there will be no NET change It will not look like there are changes taking place!

17 Evaporation is a cooling process Cooling occurs because particles with the highest energy escape first Particles left behind have lower average kinetic energies; thus the temperature decreases Similar to removing the fastest runner from a race- the remaining runners have a lower average speed Evaporation helps to keep our skin cooler on a hot day!

18 The Nature of Liquids A liquid will evaporate faster when heated because the added heat increases the average kinetic energy needed to overcome the attractive forces so more particles have enough energy to ‘escape’!

19 If you were to add a drop of water below the tube to the left what would happen? It would rise to the top & evaporate. What would it do to the surface of the mercury? (push it down!) Vapor Pressure – pressure exerted by vapor!

20 ( Since different liquids evaporate at different rates, they have different vapor pressures (at the same temps)

21 ( Which liquid is the most volatile?
Yes! Diethyl Ether! It depressed the mercury the most (highest vapor pressure!)

22 Vapor Pressures of Liquids
Which is the most volatile liquid here? Diethyl ether It has the highest vapor pressure at any Temp. Which has the strongest forces of attraction? Water

23 A liquid boils when its vapor pressure equals the external pressure, so the boiling point changes if the external pressure changes. Bubbles form throughout the liquid, rise to the surface, and escape into the air Normal boiling point- is when the vapor pressure of a liquid equals standard pressure. (1 atm)

24 The boiling point (bp) is the temperature at which the vapor pressure of the liquid is equal to the external pressure on the liquid VP BP Direct relationship!

25 Normal bp of water = 100 oC However, in Denver = 95 oC, since Denver is 1600 m above sea level and average atmospheric pressure is about 85.3 kPa (Recipe adjustments?) In pressure cookers, which reduce cooking time, water boils above 100 oC due to the increased pressure

26 Boiling Point of Water at Various Locations
Feet above sea level Patm (kPa) Boiling Point (C) Top of Mt. Everest, Tibet 29,028 32 70 Top of Mt. Denali, Alaska 20,320 45.3 79 Top of Mt. Whitney, California 14,494 57.3 85 Top of Mt. Washington, N.H. 6,293 78.6 93 Boulder, Colorado 5,430 81.3 94 New York City, New York 10 101.3 100 Death Valley, California -282 102.6 100.3

27 Vapor Pressures of Liquids
Normal bp when crossing here At any pt. on a curve line, liquid is boiling

28 SOLIDS If you cool a liquid, the particles lose kinetic energy and slow down. If they slow down enough, extra forces of attraction pull them in so close together that they can only vibrate in place. Freezing – change of a liquid to a solid.

29 Types of Solids Molecular solids Metallic solids Ionic solids
Covalent network solids (diamonds)

30 Crystals or Crystalline Solids
Particles of crystals are arranged in repeating geometric patterns NaCl

31 Representation of Components in a Crystalline Solid
Crystal Lattice: The orderly, regular, 3-dimensional arrangement of particles (atoms, ions, etc.) in a crystal.

32 Unit Cell The smallest portion of a crystal lattice that shows the three-dimensional pattern of the entire lattice Ex. A cubic lattice system has 3 types of unit cells

33 Table salt crystals are shaped like cubes.

34 Diamond, a form of carbon, is also a crystalline solid.
the crystals are shaped something like pyramids.

35 Non-crystalline solids
Many solids do not form crystals- Amorphous Their molecules do not arrange into repeating patterns often because they are too large. No definite melting point Examples: Glass - also called a super-cooled liquid many plastics, soot, asphalt, butter

36 PHASE CHANGES PHASE CHANGES – change is physical state (melting, freezing, boiling, condensing, sublimation, deposition) BOTH PHASES present during a phase change Temperature remains constant during a phase change. Sublimation – change of a solid directly to a gas (dry ice, iodine, snow) Deposition – change of a gas directly to a solid.

37 Temperature (C°) 60° 20° -20° Heat (kilojoules)

38 Heating and cooling curve for water heated at a constant rates.
A-B = Solid ice, temperature is increasing. Particles gain kinetic energy, vibration of particles increases. Ice

39 B-C = Solid starts to change state from solid to liquid
B-C = Solid starts to change state from solid to liquid. Temperature remains constant as energy is used to break inter-molecular bonds. H2O (s)  H2O (l) 0ºC

40 C-D = temperature starts to rise once all the solid has melted
C-D = temperature starts to rise once all the solid has melted. Particles gain kinetic energy. Liquid water

41 D-E = Liquid starts to vaporize, turning from liquid to gas
D-E = Liquid starts to vaporize, turning from liquid to gas. The temperature remains constant as energy is used to break inter-molecular forces. H2O ()  H2O (g) 100ºC

42 E-F = temperature starts to rise once all liquid is vaporized
E-F = temperature starts to rise once all liquid is vaporized. Gas particles gain kinetic energy. steam

43 The heating/cooling curve for water

44 Water phase changes constant Temperature remains __________
during a phase change. constant Water phase changes KE not changing during phase changes Boiling condensation Melting freezing Kinetic energy increasing on slopes

45 Phase Diagram Represents phases as a function of temperature and pressure.

46 Classification of Matter
Recall Classification of Matter Matter Pure Substances Elements Compounds Mixtures Homogeneous Heterogeneous Also called solutions

47 Solute Solvent Aqueous Solutions – water solutions
The part of a solution that gets dissolved – the part in lesser quantity Salt in salt water Sugar in soda drinks Carbon dioxide in soda drinks Solvent The part of a solution that does the dissolving – the part in greater quantity Water in salt water Water in soda

48 Solutions occur in all 3 phases!
Solute Solvent Example Gas Air Liquid Water Vapor in Air Solid Mothballs Seltzer Water Antifreeze in radiator, cocktail Salt water Whipped cream Fillings: Hg in Ag Alloys: Brass, etc.

49 non-polar molecules are soluble in non-polar solvents
Two substances with similar intermolecular forces are likely to be soluble in each other. non-polar molecules are soluble in non-polar solvents Ex. Grease in gasoline Ionic compounds & polar molecules are soluble in polar solvents Ex. Ethanol in water; salt in water “like dissolves like” 12.2

50 When a salt dissolves in water, the positive
ions are attracted to the slightly negative ends of the water molecules and the negative ions are attracted to the slightly positive ends of the water molecules. They ‘dissociate’ (separate) The ions are more strongly attracted to each other. but they become surrounded by water molecules and can’t get back together! Solvation – where solvent molecules surround solute particles.

51 Dissolving of solid sodium chloride.

52 Water: the Universal Solvent
Cl- Na+ O O O


54 SOLUBILITY A measure of how much a gas, liquid, or solid will dissolve
in a solvent.

55 Factors Effecting Solubility
The solubility of MOST solids increases with an increase in temperature. The rate at which solids dissolve increases with increasing surface area of the solid. (crush them & stir them!) The solubility of gases increases with decreases in temperature. The solubility of gases increases with an increase of pressure above the solution.

56 Therefore… Solids tend to dissolve best when: Heated Stirred
Ground into smaller particles Gases tend to dissolve best when: The solution is cooler Pressure is higher

57 When you see bubbles, it’s a heterogeneous mixture not a solution!
CO2 in and out of water Mixture Solution When you see bubbles, it’s a heterogeneous mixture not a solution!

58 Salt (NaCl) is very soluble in water. 350 g/liter
Salt (NaCl) is very soluble in water g/liter. However, if water evaporates, there will be too much salt for the water to hold in solution. The salt begins to form crystals. A lake near Death Valley is supersaturated with salt causing the salt to crystallize out.

59 Once a year the company who owns the lakes lets visitors into the area to collect salt crystals.

60 These are some of the salt crystals collected.

61 Solubility maximum mass of solute that will dissolve in 100 g of solvent at a given temperature varies with temp based on a saturated solution

62 Temperature and Solubility
Solid solubility and temperature solubility increases with increasing temperature solubility decreases with increasing temperature 12.4

63 Solubility Chart

64 A solution that contains the maximum amount of solute that may be dissolved under existing conditions is saturated. A solution that contains less solute than it can is unsaturated. A solution that contains more dissolved solute than it should is supersaturated.

65 Solubility UNSATURATED SOLUTION more solute dissolves
no more solute dissolves SUPERSATURATED SOLUTION becomes unstable, crystals form concentration

66 Solubility Chart

67 Colligative Properties
Colligative Properties are those that depend on the concentration of particles in a solution, not upon the identity of the solute. Boiling Point Elevation Freezing Point Depression Vapor Pressure Depression More particles in solution – bigger effect on the coll. property! Higher bp, lower fp, etc.

68 Change in Boiling Point
Common Applications of Boiling Point Elevation

69 Change in Freezing Point
Common Applications of Freezing Point Depression Ethylene glycol – deadly to small animals Propylene glycol

70 Concentration

71 Concentrations of Solutions
The amount of solute in a solution is given by its concentration. Concentration can be expressed qualitatively or quantitatively – Dilute – contains small amount of solute Concentrated – contains large amount of solute (words) (numbers)

72 Molarity = ( M ) moles solute liters of solution
We will use something called Molarity! Molarity ( M ) = moles solute liters of solution *Volume of solution MUST be in Liters! Ex. What is the molarity of a solution made by dissolving 3.5 moles of NaCl in a 4.3L solution? 3.5 moles NaCl = 0.81 M NaCl(aq) 4.3 L solution Molarity can be labeled “M” or “mol/L” Handy in conversion factors!

73 PROBLEM: Dissolve 5.00 g of NiCl2 in enough water to make 250 mL of solution. Calculate the Molarity. Step 1: Calculate moles of NiCl2 Step 2: convert mL to L 250 mL L Step 3: Calculate Molarity [NiCl2] = M

74 moles = M•V What mass of oxalic acid, H2C2O4, is
required to make 250. mL of a M solution? *Rearrange formula to solve for moles. moles = M•V Step 1: Change mL to L. 250 mL x 1L/1000mL = L Step 2: Calculate. Moles = ( mol) (0.250 L) = moles Step 3: Convert moles to grams. ( mol)(90.00 g/mol) = g oxalic acid 1 L

75 Learning Check How many grams of NaOH are required to prepare 400. mL of 3.0 M NaOH solution? 1) 12 g 2) 48 g 3) g

76 Preparing Solutions M1V1 = M2V2
Weigh out a specific mass of solid solute and dissolve in a given quantity of solvent. Dilute a concentrated solution to give one that is less concentrated. M1V1 = M2V2

77 How would you make 250 mL of 3.0 M HCl from a 6 M stock solution of HCl? M1V1 = M2V2 (250 mL) (3.0 M) = (6 M) (X) X = 125 mL “How would you make…” would require some words in the answer! So,… Add enough water to 125 mL of 6 M HCl to bring the total volume to 250 mL.

78 Miscible – where liquids dissolve in each other.
Ex. Ethanol in water

79 Suspensions and Colloids
Suspensions and colloids are NOT solutions. Suspensions: The particles are so large that they settle out of the solvent if not constantly stirred. Colloids: The particles are intermediate in size between those of a suspension and those of a solution.

80 The Tyndall Effect Colloids scatter light, making a beam visible. Solutions do not scatter light. Which glass contains a colloid? colloid solution

81 Types of Colloids Examples Dispersing Medium Dispersed Substance
Colloid Type Fog, aerosol sprays Gas Liquid Aerosol Smoke, airborne bacteria Solid Whipped cream, soap suds Foam Milk, mayonnaise Emulsion Paint, clays, gelatin Sol Marshmallow, Styrofoam Solid foam Butter, cheese Solid emulsion Ruby glass Solid sol

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