2 13.1 The Solution ProcessSolutions are homogeneous mixtures of two or more pure substances.Formed when 1 substance (solute) disperses uniformly throughout another (solvent).Solute – present in lesser quantitySolvent – present in greater quantity
3 Solutions can be mixes involving any state. 13.1 The Solution ProcessSolutions can be mixes involving any state.This mix always forms a solution!
4 How Does a Solution Form? Solvation As a solution forms, the solvent pulls solute particles apart and surrounds, or solvates, them.If the solvent is water, this process is termed hydration.
5 How Does a Solution Form? Dispersion In the absence of a strong solvent-solute attraction, the particles randomly spread out or Disperse.This is typical for nonpolar solutions!
6 Why do Solutions Form?The ability of substances to form solutions depends on 2 factors:Type of intermolecular interactions involvedNatural tendency of substances to spread into larger volumes when not restrained
7 The Effect of Intermolecular Forces There are 3 forces at work during solvation.Break Solvent-Solvent attraction.Break solute-solute attraction.Make solvent-solute attractionIf #3 is comparable to, or greater, than #1 and #2, than a solutions forms!
8 Break Solvent – Solvent Forces H to N,O, or FHydrogenbondingPolarSolventDipole-dipoleNonpolarsolventLondonorDispersion(induced dipoles)
9 Break Solute – Solute Forces H to N,O, or FHydrogenbondingDipole-dipolePolar MoleculesLondonorDispersion(induced dipoles)NonpolarMolecules
11 NaCl and WaterNaCl is soluble in water, it is because the ion-dipole interactions are strong enough to overcome the lattice energy of the salt crystal.
12 The solute-solvent force is too weak to overcome the ionic bonding! NaCl and C6H14+>Solute-solventInduced dipole(weak)Solute-soluteionic(strong)Solvent-solventLondon dispersion(weak)Insoluble!The solute-solvent force is too weak to overcome the ionic bonding!
13 Energy Changes and Solution Formation Simply put, three processes affect the energetics of solution:separation of solute particles,endothermicseparation of solvent particles,new interactions between solute and solvent.exothermic
14 Energy Changes in Solution The enthalpy change of the overall process depends on H for each of these steps.
15 Energy Changes in Solution Things do not tend to occur spontaneously (i.e., without outside intervention) unless the energy of the system is lowered.(overall exothermic)
16 Why Do Endothermic Processes Occur? Yet we know that in some processes, like the dissolution of NH4NO3 in water, heat is absorbed, not released.
17 Solution Formation, Spontaneity, and Entropy This is where the second factor effecting solutions is involves!Natural tendency of substances to spread into larger volumes when not restrained
18 Solution Formation, Spontaneity, and Entropy Dissolving, when it occurs, is a spontaneous process!If you remember from chemistry, reactions that increase the entropy (disorder and randomness) tend to be spontaneous.
19 Enthalpy Is Only Part of the Picture The reason is that increasing the disorder or randomness (known as entropy) of a system tends to lower the energy of the system.Diffusion is therefore a spontaneous reaction
20 Enthalpy Is Only Part of the Picture So even though enthalpy may increase (endothermic), the overall energy of the system can still decrease if the system becomes more disordered.
21 Silver chloride is essentially insoluble in water. Would you expect a significant change in entropy of the system when 10 g of AgCl is added to 100 g of water?NO!
22 Solution Formation and Chemical Reactions Just because a substance disappears when it comes in contact with a solvent, it doesn’t mean the substance dissolved.
23 Solution Formation and Chemical Reactions Dissolution is a physical change — you can get back the original solute by evaporating the solvent.If you can’t, the substance didn’t dissolve, it reacted.
24 13.2 Types of Solutions Saturated In a saturated solution, the solvent holds as much solute as is possible at that temperature.Dissolved solute is in dynamic equilibrium with solid solute particles.
25 13.2 Types of Solutions Unsaturated If a solution is unsaturated, less solute than can dissolve in the solvent at that temperature is dissolved in the solvent.
26 13.2 Types of Solutions Supersaturated In supersaturated solutions, the solvent holds more solute than is normally possible at that temperature.These solutions are unstable; crystallization can usually be stimulated by adding a “seed crystal” or scratching the side of the flask.
27 13.3 Factors Affecting Solubility Chemists use the axiom “like dissolves like."Polar substances tend to dissolve in polar solvents.Nonpolar substances tend to dissolve in nonpolar solvents.
28 Solute-Solvent Interactions The more similar the intermolecular attractions, the more likely one substance is to be soluble in another.
29 Solute-Solvent Interactions Glucose (which has hydrogen bonding) is very soluble in water, while cyclohexane (which only has dispersion forces) is not.
30 Decrease, -CH3 is nonpolar What would happen to the solubility of glucose if the hydroxyl groups (-OH) were replaced with methyl groups (-CH3)?Decrease, -CH3 is nonpolar
31 Solute-Solvent Interactions Vitamin A is soluble in nonpolar compounds (like fats).Vitamin C is soluble in water.
32 Pressure EffectsIn general, the solubility of gases in water increases with increasing mass.Larger molecules have stronger dispersion forces.
33 Pressure EffectsThe solubility of liquids and solids does not change appreciably with pressure.The solubility of a gas in a liquid is directly proportional to its pressure.
34 Henry’s Law Sg = kPg where Sg is the solubility of the gas, k is the Henry’s Law constant for that gas in that solvent, andPg is the partial pressure of the gas above the liquid.
35 Henry’s LawHow can this relationship be expressed?
36 Henry’s Law Problem x = (4.0 atm)(0.030 M)/(1.0 atm) x = 0.12 M Calculate the concentration of CO2 in a soft drink that is bottled with a partial pressure of CO2 at atm over the liquid at 25 ºC if the solubility at 1.0 atm is M?x = (4.0 atm)(0.030 M)/(1.0 atm)x = 0.12 M
37 Temperature EffectsGenerally, the solubility of solid solutes in liquid solvents increases with increasing temperature.
38 Why do bubbles form on the inside wall of a cooking pot when the water is heated on the stove, even though the temperature is well below the boiling point of water?Increasing the temperature of the water decreases the solubility of the dissolved gasses. These released gasses deposit on the side of the container.
39 Temperature Effects The opposite is true of gases. Carbonated soft drinks are more “bubbly” if stored in the refrigerator.Warm lakes have less O2 dissolved in them than cool lakes.
40 13.4 Ways of Expressing Concentrations Qualitative ConcentrationsDiluteRelatively small concentrationConcentratedRelatively large concentrationQuantitative ConcentrationMass percentppm and ppbMole fractionMolarityMolality
41 % mass is also known as “pph” (parts per hundred) Mass Percentagemass of A in solutiontotal mass of solution 100What does a 4.0 % NaCl solution mean?4.0 g NaCl for every 96.0 g of water!% mass is also known as “pph” (parts per hundred)
42 Parts per Million and Parts per Billion Parts per Million (ppm)mass of A in solutiontotal mass of solutionppm = 106Parts per Billion (ppb)mass of A in solutiontotal mass of solutionppb = 109
43 total moles in solution Mole Fraction (χ)moles of Atotal moles in solutionXA =In some applications, one needs the mole fraction of solvent, not solute — make sure you find the quantity you need!Xsolute + Xsolvent = 1
44 Molarity (M) mol of solute M = L of solution You will recall this concentration measure from Chapter 4.Since volume is temperature-dependent, molarity can change with temperature.
45 Molality (m) mol of solute m = kg of solvent Since both moles and mass do not change with temperature, molality (unlike molarity) is not temperature-dependent.
46 Changing Molarity to Molality If we know the density of the solution, we can calculate the molality from the molarity and vice versa. (p545)
47 13.5 Colligative Properties Changes in colligative properties depend only on the number of solute particles present, not on the identity of the solute particles.Among colligative properties areVapor pressure loweringBoiling point elevationMelting point depressionOsmotic pressure
48 Vapor PressureBecause of solute-solvent intermolecular attraction, higher concentrations of nonvolatile solutes make it harder for solvent to escape to the vapor phase.
49 Vapor PressureTherefore, the vapor pressure of a solution is lower than that of the pure solvent.
50 Raoult’s Law PA = APA where A is the mole fraction of compound A, andPA is the normal vapor pressure of A at that temperature.NOTE: This is one of those times when you want to make sure you have the vapor pressure of the solvent.
51 Boiling Point Elevation and Freezing Point Depression Nonvolatile solute-solvent interactions also cause solutions to have higher boiling points and lower freezing points than the pure solvent.
52 Boiling Point Elevation The change in boiling point is proportional to the molality of the solution:Tb = Kb mwhere Kb is the molal boiling point elevation constant, a property of the solvent.Tb is added to the boiling point of the pure solvent.
53 Freezing Point Depression The change in freezing point can be found similarly:Tf = Kf mHere Kf is the molal freezing point depression constant of the solvent.Tf is subtracted from the boiling point of the pure solvent.
54 Boiling Point Elevation and Freezing Point Depression Note that in both equations, T does not depend on what the solute is, but only on how many particles are dissolved.Tb = Kb mTf = Kf m
55 Colligative Properties of Electrolytes Since these properties depend on the number of particles dissolved, solutions of electrolytes (which dissociate in solution) should show greater changes than those of nonelectrolytes.
56 van’t Hoff FactorWe use the van’t Hoff factor, i, to represent the number of particles we get when a substance dissolves.Nonelectrolytes (molecules)i = 1Electrolytes (ionic compounds)i = # of ionsNaCl → i = 2K2SO4 → i = 3MgSO4 → i = 2
57 van’t Hoff Factor Tf = Kf m i Tb = Kb m i We modify the previous equations by multiplying by the van’t Hoff factor, i.Tf = Kf m iTb = Kb m i
58 van’t Hoff FactorHowever, a 1M solution of NaCl does not show twice the change in freezing point that a 1M solution of methanol does.
59 van’t Hoff FactorOne mole of NaCl in water does not really give rise to two moles of ions.
60 van’t Hoff FactorSome Na+ and Cl- reassociate for a short time, so the true concentration of particles is somewhat less than two times the concentration of NaCl.
63 OsmosisSome substances form semipermeable membranes, allowing some smaller particles to pass through, but blocking other larger particles.In biological systems, most semipermeable membranes allow water to pass through, but solutes are not free to do so.
64 OsmosisIn osmosis, there is net movement of solvent from the area of higher solvent concentration (lower solute concentration) to the are of lower solvent concentration (higher solute concentration).
65 Osmotic Pressure n = ( )RT = MRT V The pressure required to stop osmosis, known as osmotic pressure, , isnV = ( )RT = MRTwhere M is the molarity of the solution.If the osmotic pressure is the same on both sides of a membrane (i.e., the concentrations are the same), the solutions are isotonic.
66 Osmosis in Blood CellsIf the solute concentration outside the cell is greater than that inside the cell, the solution is hypertonic.Water will flow out of the cell, and crenation results.
67 Osmosis in CellsIf the solute concentration outside the cell is less than that inside the cell, the solution is hypotonic.Water will flow into the cell, and hemolysis results.
68 13.6 ColloidsSuspensions of particles larger than individual ions or molecules, but too small to be settled out by gravity are called colloids.
69 Tyndall Effect Colloidal suspensions can scatter rays of light. This phenomenon is known as the Tyndall effect.
70 Mixtures Review Settle out? Particle size Tyndall Effect Type of mixtureSuspensionColloidSolutionyesyeslargeheterogeneousnointermediateyesheterogeneousnosmallnohomogeneous
71 Colloids in Biological Systems Some molecules have a polar, hydrophilic (water-loving) end and a non-polar, hydrophobic (water-hating) end.
72 Hydrophilic ColloidIn this molecules the hydrophilic parts are on the surface end and the hydrophobic parts are folded in to make a hydrophilic colloid.Hydrophilic parts become attracted to water and partially hydrated.72
73 Hydrophilic ColloidIn this molecules the hydrophilic parts are on the surface end and the hydrophobic parts are folded in to make a hydrophilic colloid.Hydrophilic parts become attracted to water and partially hydrated.73
74 How do hydrophilic colloid form? If ions get adsorbed (stick to surface), it causes them to be attracted by water.Ions stabilize the colloid and allow limited hydration and prevents clumping nonpolar colloid by ion to ion repulsion.
75 How do hydrophilic colloid form? They can also be stabilized as a colloid by emulsifying agents (molecule with polar and nonpolar portions like soap)