11 Chapter 13 Properties of Solutions CHEMISTRY The Central Science 9th Edition

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Chapter 13Chapter 13Properties of SolutionsProperties of Solutions

CHEMISTRY The Central Science

9th Edition

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Text, P. 417, review (Chapter 11)

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• Solutions• homogeneous mixtures

• Solution formation is affected by• strength and type of intermolecular forces • forces are between and among the solute and solvent

particles

13.1: The Solution Process13.1: The Solution Process

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Text, P. 486

Hydration of solute

• Attractive forces between solute & solvent particles are comparable in magnitude with those between the solute or solvent particles themselves

• Note attraction of charges

•What has to happen to:

• Water’s H-bonds?

• NaCl?

•What intermolecular

force is at work in

solvation?

Text, P. 486

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Energy Changes and Solution Formation

There are three energy steps in forming a solution:

• the enthalpy change in the solution process isHsoln = H1 + H2 + H3

• Hsoln can either be + or - depending on the intermolecular forces

Text, P. 487

Text, P. 488

MgSO4 Hot Pack NH4NO3 Cold Pack

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• Breaking attractive intermolecular forces is always endothermic

• Forming attractive intermolecular forces is always exothermic

• To determine whether Hsoln is positive or negative, consider the strengths of all solute-solute and solute-solvent interactions:

• H1 and H2 are both positive

H3 is always negative

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• Rule: Polar solvents dissolve polar solutes

Non-polar solvents dissolve non-polar solutes

(like dissolves like)

WHY?

– If Hsoln is too endothermic a solution will not form

– NaCl in gasoline: weak ion-dipole forces (gasoline is non-polar)

– The ion-dipole forces do not compensate for the separation of ions

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Solution Formation, Spontaneity, and Disorder

• A spontaneous process occurs without outside intervention

• When energy of the system decreases, the process is spontaneous• Some spontaneous processes do not involve the system

moving to a lower energy state (e.g. an endothermic reaction)

• If the process leads to a greater state of disorder, then the process is spontaneous• Entropy

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Example: a mixture of CCl4 and C6H14 is less ordered than the two separate liquids

•Therefore, they spontaneously mix even though Hsoln is very close to zero

Text, P. 489

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Solution Formation and Chemical Reactions• Example:

Ni(s) + 2HCl(aq) NiCl2(aq) + H2(g)

• When all the water is removed from the NiCl2 solution, no Ni is found only NiCl2·6H2O (a chemical reaction that results

in the formation of a solution)

• Water molecules fit into the crystal lattice in places not specifically occupied by a cation or an anion

• Hydrates• Water of hydration

• Think about it: What happens when NaCl is dissolved in water and then heated to dryness?

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NaCl(s) + H2O (l) Na+(aq) + Cl-(aq)

• When the water is removed from the solution, NaCl is found• NaCl dissolution is a physical process

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• Sample problem # 3

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• Dissolve: solute + solvent solution• Crystallization: solution solute + solvent• Saturation: crystallization and dissolution are in

equilibrium• Solubility: amount of solute required to form a saturated

solution• Supersaturated: a solution formed when more solute is

dissolved than in a saturated solution

13.2: Saturated Solutions and 13.2: Saturated Solutions and SolubilitySolubility

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1. Solute-Solvent Interaction• “Like dissolves like”• Miscible liquids: mix in any proportions• Immiscible liquids: do not mix

13.3: Factors Affecting 13.3: Factors Affecting SolubilitySolubility

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Generalizations:

• Intermolecular forces are important: • Water and ethanol are miscible

• broken hydrogen bonds in both pure liquids are

re-established in the mixture

• The number of carbon atoms in a chain affects solubility: the more C atoms in the chain, the less soluble the substance is in water

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Generalizations, continued:

• The number of -OH groups within a molecule increases solubility in water

• The more polar bonds in the molecule, the better it dissolves in a polar solvent (like dissolves like)

• Network solids do not dissolve• the strong IMFs in the solid are not re-established in any

solution

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Text, P. 493

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Fat soluble vitamin Water soluble

vitamin

Read “Chemistry & Life”, P. 494

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2. Pressure Effects• Solubility of a gas in a liquid is a function of the pressure

of the gas

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• High pressure means • More molecules of gas are close to the solvent• Greater solution/gas interactions• Greater solubility

• If Sg is the solubility of a gas

k is a constant

Pg is the partial pressure of a gas

then Henry’s Law gives:

Carbonated Beverages!

gg kPS

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3. Temperature Effects

• As temperature increases• Solubility of solids

generally increases• Solubility of gases

decreases• Thermal pollution

Text, P. 497

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Figure 13.17, P. 497

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• Sample problem # 17

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• All methods involve quantifying amount of solute per amount of solvent (or solution)• Amounts or measures are masses, moles or liters• Qualitatively solutions are dilute or concentrated

13.4: Ways of Expressing 13.4: Ways of Expressing ConcentrationConcentration

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610solution of mass total

solutionin component of masscomponent of ppm

910solution of mass total

solutionin component of masscomponent of ppb

100solution of mass total

solutionin component of masscomponent of % mass

• Definitions:

1.

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2.

3.

• Recall mass can be converted to moles using the molar mass

solution of moles total

solutionin component of molescomponent offraction Mole

solution of literssolute moles

Molarity

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4.

• Converting between molarity (M) and molality (m) requires density• Molality doesn’t vary with temperature

• Mass is constant• Molarity changes with temperature

• Expansion/contraction of solution changes volume

solvent of kgsolute moles

Molality, m

Text, P. 501

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• Sample Problems #31, 33, 37, 39, 41

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Colligative properties depend on quantity of solute particles, not their identity• Electrolytes vs. nonelectrolytes

0.15m NaCl 0.15m in Na+ & 0.15m in Cl- 0.30m in particles

0.050m CaCl2 0.050m in Ca+2 & 0.1m in Cl- 0.15m in particles

0.10m HCl 0.10m in H+ & 0.10m in Cl- 0.20m in particles

0.050m HC2H3O2 between 0.050m & 0.10m in particles

0.10m C12H22O11 0.10m in particles

• Compare physical properties of the solution with those of the pure solvent

13.5: Colligative Properties13.5: Colligative Properties

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1. Lowering Vapor Pressure

• Non-volatile solutes reduce the ability of the surface solvent molecules to escape the liquid

• Vapor pressure is lowered

• Raoult’s Law:

PA is the vapor pressure with solute

PA is the vapor pressure without solute

A is the mole fraction of solvent in solution A

AAA PP

Increase X of solute, decrease vapor pressure above the solution

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Ideal solution: one that obeys Raoult’s law• Raoult’s law breaks down (Real solutions)

• Real solutions approximate ideal behavior when • solute concentration is low• solute and solvent have similar IMFs

• Assume ideal solutions for problem solving

2. Boiling-Point Elevation• The triple point - critical point curve is lowered

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• At 1 atm (normal BP of pure liquid) there is a lower vapor pressure of the solution• A higher temperature is required to reach a vapor

pressure of 1 atm for the solution (Tb)

• Molal boiling-point-elevation constant, Kb, expresses how much Tb changes with molality, m:

mKT bb

Text, P. 505

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3. Freezing Point Depression

• The solution freezes at a lower temperature (Tf) than the pure solvent– lower vapor pressure for the solution

• Decrease in FP (Tf) is directly proportional to molality (Kf is the molal freezing-point-depression constant):

mKT ff

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Text, P. 505

Applications: Antifreeze!

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• Examples: # 45, 47, 49, 51 & 53

• A neat link

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4. Osmosis

• Semipermeable membrane: permits passage of some components of a solution• Example: cell membranes and cellophane

• Osmosis: the movement of a solvent from low solute concentration to high solute concentration• There is movement in both directions across a

semipermeable membrane• “Where ions go, water will flow” ~ Mrs. Moss

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• Eventually the pressure difference between the arms stops osmosis

Text, P. 507

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• Osmotic pressure, , is the pressure required to stop osmosis:

• It is colligative because it depends on the concentration of the solute in the solvent

MRT

RTVn

nRTV

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• Isotonic solutions: two solutions with the same separated by a semipermeable membrane

• Hypertonic solution: a solution that is more concentrated than a comparable solution

• Hypotonic solution: a solution of lower than a hypertonic solution

• Osmosis is spontaneous• Read text, P. 508 – 509 for practical examples

461

• Examples: #57, 59 & 61

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• There are differences between expected and observed changes due to colligative properties of strong electrolytes

– Electrostatic attractions between ions

– “ion pair” formation temporarily reduces the number of particles in solution

– van’t Hoff factor (i): measure of the extent of ion dissociation

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• Ratio of the actual value of a colligative property to the calculated value (assuming it to be a nonelectrolyte)– Ideal value for a salt is the # of ions per formula unit

rolyte)f(nonelect

)f(measured

T

Ti

Factors that affect i:

•Dilution

•Magnitude of charge on ions

• lower charges, less deviation

491

• Sample Problem, # 63, 82

501

• Read Text, Section 13.6, P. 511 – 515– Terms/Processes:

• Tyndall effect

• Hydrophilic

• Hydrophobic

• Adsorption

• Coagulation

11.6: Colloids11.6: Colloids

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• Read Text, Section 13.6, P. 511 – 515• Suspensions in which the suspended particles are larger than

molecules• too small to drop out of the suspension due to gravity

• Tyndall effect: ability of a colloid to scatter light• The beam of light can be seen through the colloid

11.6: Colloids11.6: Colloids

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Text, P. 512

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Hydrophilic and Hydrophobic Colloids

• “Water loving” colloids: hydrophilic• “Water hating” colloids: hydrophobic

• Molecules arrange themselves so that hydrophobic portions are oriented towards each other

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• Adsorption: when something sticks to a surface we say that it is adsorbed• Ions stick to a colloid (colloids appears hydrophilic)

• Oil drop and soap (sodium stearate)• Sodium stearate has a long hydrophobic tail (Carbons)

and a small hydrophilic head (-CO2-Na+)

Text, P. 514

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Removal of Colloidal Particles

• Coagulation (enlarged) until they can be removed by filtration

• Methods of coagulation:– heating (colloid particles are attracted to each other when they

collide)

– adding an electrolyte (neutralize the surface charges on the colloid particles)

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End of Chapter 13End of Chapter 13Properties of SolutionsProperties of Solutions