2 Intermolecular Forces Forces of attraction between different molecules rather than bonding forces within the same molecule.Dipole-dipole attractionBetween particles with charged sidesmlHydrogen bonds - StrongestBetween H’s & F, O, N on other moleculesDispersion forces - WeakestCaused 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 sharedelectrons longer than the hydrogen atoms. As aresult the oxygen side isnegatively charged and the hydrogen side of water is positively charged.O
4 It is its polarity andhydrogen bonding that givewater its many unusualProperties!i.e. high boiling point, expansion upon freezing, and surface tension
5 Hydrogen bonds - Strongest Between H’s & F, O, N on other molecules IceLiquid
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 itselfso 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.OOOil dropletOOO
8 Forces and PhasesSubstances with very little intermolecular attraction exist as gasesSubstances with strong intermolecular attraction exist as liquidsSubstances with very strong intermolecular (or ionic) attraction exist as solids
9 “V” is for Vocabulary!Vapor – gaseous state of asubstance that is not normallya gas at room temperature.Volatile – evaporates rapidly(due to weak intermolecular forces)Ex. ‘thin’ liquidsViscous – evaporates slowly(due to strong intermolecular forces)Ex. ‘thick’ liquids
12 LiquidsA 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 gasto a liquidIt is now inliquid form!
13 The Nature of LiquidsThe 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 EvaporationParticles near the surface with enough kinetic energy that happen to bounce in the right direction escape!
15 The Nature of LiquidsEventually 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 condensationNote that there will still be particles that evaporate and condenseBut, there will be no NET changeIt will not look like there are changes taking place!
17 Evaporation is a cooling process Cooling occurs because particles with the highest energy escape firstParticles left behind have lower average kinetic energies; thus the temperature decreasesSimilar to removing the fastest runner from a race- the remaining runners have a lower average speedEvaporation helps to keep our skin cooler on a hot day!
18 The Nature of LiquidsA 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 adrop of water belowthe tube to the leftwhat would happen?It would rise to thetop & evaporate.What would it do tothe surface of themercury?(push it down!)Vapor Pressure –pressure exerted byvapor!
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 etherIt has the highest vapor pressure at anyTemp.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 airNormal 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 liquidVPBPDirect relationship!
25 Normal bp of water = 100 oCHowever, 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 levelPatm (kPa)Boiling Point (C)Top of Mt. Everest, Tibet29,0283270Top of Mt. Denali, Alaska20,32045.379Top of Mt. Whitney, California14,49457.385Top of Mt. Washington, N.H.6,29378.693Boulder, Colorado5,43081.394New York City, New York10101.3100Death Valley, California-282102.6100.3
27 Vapor Pressures of Liquids Normal bp whencrossing hereAt any pt. on acurve line,liquid is boiling
28 SOLIDSIf 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.
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- AmorphousTheir molecules do not arrange into repeating patternsoften because they are too large.No definite melting pointExamples:Glass - also called a super-cooled liquidmany plastics, soot, asphalt, butter
36 PHASE CHANGESPHASE CHANGES – change is physical state (melting, freezing, boiling, condensing, sublimation, deposition)BOTH PHASES present during a phase changeTemperature 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.
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
44 Water phase changes constant Temperature remains __________ during a phase change.constantWater phase changesKE not changing duringphase changesBoilingcondensationMeltingfreezingKinetic energy increasing on slopes
45 Phase DiagramRepresents phases as a function of temperature and pressure.
46 Classification of Matter RecallClassification of MatterMatterPureSubstancesElementsCompoundsMixturesHomogeneousHeterogeneousAlso called solutions
47 Solute Solvent Aqueous Solutions – water solutions The part of a solution that gets dissolved – the part in lesser quantitySalt in salt waterSugar in soda drinksCarbon dioxide in soda drinksSolventThe part of a solution that does the dissolving – the part in greater quantityWater in salt waterWater in soda
48 Solutions occur in all 3 phases! SoluteSolventExampleGasAirLiquidWater Vapor in AirSolidMothballsSeltzer WaterAntifreeze in radiator, cocktailSalt waterWhipped creamFillings: Hg in AgAlloys: 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 solventsEx. Grease in gasolineIonic compounds & polar molecules are soluble in polar solventsEx. 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 negativeends 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 eachother. but they become surrounded by water molecules and can’t get back together!Solvation – where solvent moleculessurround solute particles.
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 particlesGases tend to dissolve best when:The solution is coolerPressure is higher
57 When you see bubbles, it’s a heterogeneous mixture not a solution! CO2 in and out of waterMixtureSolutionWhen 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.
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 dissolvesSUPERSATURATED SOLUTIONbecomes unstable, crystals formconcentration
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 ElevationFreezing Point DepressionVapor Pressure DepressionMore particles in solution – bigger effecton 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 DepressionEthylene glycol – deadly to small animalsPropylene glycol
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 soluteConcentrated – contains large amount of solute(words)(numbers)
72 Molarity = ( M ) moles solute liters of solution We will use something called Molarity!Molarity(M)=moles soluteliters of solution*Volume of solution MUST be in Liters!Ex. What is the molarity of a solution madeby dissolving 3.5 moles of NaCl in a 4.3L solution?3.5 moles NaCl= 0.81 M NaCl(aq)4.3 L solutionMolarity 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 NiCl2Step 2: convert mL to L250 mL LStep 3: Calculate Molarity[NiCl2] = M
74 moles = M•V What mass of oxalic acid, H2C2O4, is required to make 250. mL of a Msolution?*Rearrange formulato solve for moles.moles = M•VStep 1: Change mL to L.250 mL x 1L/1000mL = LStep 2: Calculate.Moles = ( mol) (0.250 L) = molesStep 3: Convert moles to grams.( mol)(90.00 g/mol) = g oxalic acid1 L
75 Learning CheckHow many grams of NaOH are required to prepare 400. mL of 3.0 M NaOH solution?1) 12 g2) 48 g3) 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 of3.0 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 EffectColloids scatter light, making a beam visible. Solutions do not scatter light.Which glass contains a colloid?colloidsolution
81 Types of Colloids Examples Dispersing Medium Dispersed Substance Colloid TypeFog, aerosol spraysGasLiquidAerosolSmoke, airborne bacteriaSolidWhipped cream, soap sudsFoamMilk, mayonnaiseEmulsionPaint, clays, gelatinSolMarshmallow, StyrofoamSolid foamButter, cheeseSolid emulsionRuby glassSolid sol