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

Text, P. 417, review (Chapter 11)

13.1: The Solution Process Solutions homogeneous mixtures Solution formation is affected by strength and type of intermolecular forces forces are between and among the solute and solvent particles

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

Energy Changes and Solution Formation There are three energy steps in forming a solution: the enthalpy change in the solution process is Hsoln = 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

H1 and H2 are both positive 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

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

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

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

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

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

Sample problem # 3

13.2: Saturated Solutions and Solubility 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

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

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

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

Text, P. 493

Read “Chemistry & Life”, P. 494 Fat soluble vitamin Water soluble vitamin

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

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!

As temperature increases Solubility of solids generally increases 3. Temperature Effects As temperature increases Solubility of solids generally increases Solubility of gases decreases Thermal pollution Text, P. 497

Figure 13.17, P. 497

Sample problem # 17

13.4: Ways of Expressing Concentration 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

Definitions: 1.

2. 3. Recall mass can be converted to moles using the molar mass

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

Text, P. 501

Sample Problems #31, 33, 37, 39, 41

13.5: Colligative Properties 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

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 Increase X of solute, decrease vapor pressure above the solution

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

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:

Text, P. 505

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):

Applications: Antifreeze! Text, P. 505 Applications: Antifreeze!

Examples: # 45, 47, 49, 51 & 53 A neat link

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

Eventually the pressure difference between the arms stops osmosis Text, P. 507

Osmotic pressure, , is the pressure required to stop osmosis: It is colligative because it depends on the concentration of the solute in the solvent

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

Examples: #57, 59 & 61

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

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 Factors that affect i: Dilution Magnitude of charge on ions lower charges, less deviation

Sample Problem, # 63, 82

11.6: Colloids Read Text, Section 13.6, P. 511 – 515 Terms/Processes: Tyndall effect Hydrophilic Hydrophobic Adsorption Coagulation

11.6: Colloids 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

Text, P. 512

Hydrophilic and Hydrophobic Colloids “Water loving” colloids: hydrophilic “Water hating” colloids: hydrophobic Molecules arrange themselves so that hydrophobic portions are oriented towards each other

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

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)

End of Chapter 13 Properties of Solutions