C H A P T E R 1 3 . P RO P E RT I E S O F S O L U T I O N S
Student LEARNING OUTCOMES
• Understand the solvation process
• Understand relationship between intermolecular forces & solubility
• Be able to understand & describe the Enthalpy changes involved in solution process
• Be able to understand what is & the role entropy plays in solutions
• Be able to understand the difference in solutions; unsaturated, saturated, supersaturated
• Be able to recognize the rule “like dissolves like”
• Be able to describe affects of temperature on solubility
• Be able to utilize Henry’s & Raoult’s Law for partial pressure & solubility of gases
• Be able to calculate various concentration units & ppm/ppb
• Be able to understand, describe, and calculate colligative properties & outcomes, including:
BP elevation, FP depression, VP lowering,
• Be able to arrange solutions & pure solvents in appropriate order based on their: increase boil point,
decrease freezing points, solubility
• Be able to determine if a solute is solubility in a particular solvent
• Be able to: recall what is or not an electrolyte, write ions in solution, identify number of ions formed in solution
CHAPTER 13 OUTLINE
13.1 The Solution Process
• A solution: homogeneous mixture of solute & solvent
• Solutions may be gases, liquids, or solids
• Each substance present is component of solution
• solvent usually component in largest amount
• other components are solutes
The Effect of Intermolecular Forces; may need to look-up intermolecular forces (ch.11)
• NaCl (solute) dissolving in water (solvent):
• Water molecules orient on NaCl crystals
• H-bonds between water molecules have to be broken
• NaCl dissociates into Na+ and Cl–
• Ion–dipole forces form between Na+ & negative end of water dipole
• Similar ion–dipole interactions form between Cl– & positive end of water dipole
• interaction between solvent & solute called solvation
• If water is solvent, called hydration
Energy Changes and Solution Formation
• 3 steps involving energy in formation of a solution:
• 1st separation of solute molecules (∆H1),
• 2nd separation of solvent molecules (∆H2), and
• 3rd formation of solute–solvent interactions (∆H3)
• define enthalpy change in solution process as: ∆Hsoln = ∆H1 + ∆H2 + ∆H3
• ∆Hsoln either be positive or negative depending on intermolecular forces
• determine ∆Hsoln +/-, consider strengths solute–solute, solvent–solvent, and solute–solvent interactions:
• Breaking attractive intermolecular forces always endothermic: ∆H1 & ∆H2 both positive
• Forming attractive intermolecular forces always exothermic: ∆H3 is always negative
• possible to have either ∆H3 > (∆H1 + ∆H2) or ∆H3 < (∆H1 + ∆H2)
• Ex: MgSO4 added to water ∆Hsoln = –91.2 kJ/mol & NH4NO3 added to water ∆Hsoln = + 26.4 kJ/mol
• How predict if solution will form? general, solutions form if ∆Hsoln is negative
• ∆Hsoln too endothermic, solution not form
• “Rule of thumb”: polar solvents dissolve polar solutes & nonpolar solvents dissolve nonpolar solutes
• Consider mixing NaCl in gasoline: only weak interactions possible because gasoline is nonpolar
• interactions not compensate for separation of ions from one another: result, NaCl not dissolve
• Consider mixing water in octane (C8H18): water has strong H-bond
• energy required to break H-bonds not compensated by interactions between water & octane:
result, water & octane not mix
Solution Formation, Spontaneity, and Disorder
• spontaneous process occurs without outside intervention
• When energy of system decreases, process is spontaneous
• Spontaneous processes tend to be exothermic
• some spontaneous processes not involve movement of system to lower energy state, endothermic rxn
• some endothermic processes occur spontaneously
• Amount of randomness or disorder of system is its entropy
• solution formation favored by increase in entropy: mixture CCl4 & C6H14 less ordered than the 2 separate
• they spontaneously mix even though ∆Hsoln is close to zero
Solution Formation and Chemical Reactions
• Some solutions form by physical processes & some by chemical processes
• Ni(s) + 2HCl(aq)  NiCl2(aq) + H2(g) Note chemical form of subst dissolved changed (Ni  NiCl2)
• dissolution of Ni in HCl is a chemical process
+
–
• NaCl(s) + H2O (l)  Na (aq) + Cl (aq) Note water is removed from solution, NaCl is found
• NaCl dissolution is physical process: Na+(aq) + Cl–(aq)
13.2 Saturated Solutions and Solubility
• a solid dissolves, a solution forms: Solute + solvent  solution
• opposite process is crystallization: Solution  solute + solvent
• crystallization & dissolution in equilibrium with undissolved solute, solution is saturated
• no further increase in amount of dissolved solute
• Solubility is amount of solute required to form saturated solution
• solution with concentration of dissolved solute is less than solubility is unsaturated
• solution is supersaturated if more solute is dissolved than in saturated solution
13.3 Factors Affecting Solubility
• tendency of substance to dissolve in another depends on:
• nature of solute • nature of solvent • temperature
• pressure (for gases)
Solute–Solvent Interactions
• Intermolecular forces are important factor in determining solubility of solute in solvent
• stronger the attraction between solute & solvent molecules, greater the solubility
• polar liquids dissolve in polar solvents: favorable dipole–dipole interactions
• Pairs of liquids that mix in any proportions are miscible: ex: Ethanol & water are miscible liquids
• Pairs of liquids not mix are immiscible: ex: Gasoline and water are immiscible liquids
• number of carbon atoms in chain affects solubility
• greater the number of carbons in chain, more molecule behaves like a hydrocarbon
• more C atoms in the alcohol, lowers solubility in water
• Increasing number of –OH groups within molecule increases solubility in water
• Generalization: “like dissolves like”
• Substances with similar intermolecular attractive forces tend to be soluble in one another
• more polar bonds in molecule, better it dissolves in polar solvent
• less polar the molecule less likely to dissolve in polar solvent & more likely to dissolve in nonpolar solvent
• Network solids not dissolve because strong intermolecular forces in solid are not reestablished in any solution
Pressure Effects
• solubility of gas in liquid is function of pressure of gas over the solution; solubilities solids & liquids not affected by
pressure
• With higher pressure, more molecules of gas close to surface of solution & probability of molecule striking surface &
entering solution is increased, therefore, higher the pressure, greater the solubility
• solubility of gas directly proportional to partial pressure of gas above solution, called Henry's law
• Henry's law expressed as: Sg = kPg
• Sg is solubility of gas, Pg partial pressure, k = Henry’s law constant
• Henry's law constant differs for each solute–solvent pair & differs with temperature
• application of Henry's law is preparation of carbonated soda; are bottled under PCO2 > 1 atm
• bottle is opened, PCO2 decreases & solubility of CO2 decreases, therefore, bubbles of CO2 escape from solution
Temperature Effects
• Experience, sugar dissolves better in warm than cold water; as temperature increases, solubility of solids generally
increases
• Experience, carbonated beverages go flat as they get warm; gases are less soluble at higher temperatures
• environmental example is thermal pollution: lakes get too warm, CO2 & O2 less soluble, not available to plants or
animals, result can be that fish suffocate
13.4 Ways of Expressing Concentration
• methods involve quantifying amount of solute per amount of solvent (or solution)
• Concentration expressed qualitatively or quantitatively
• terms dilute & concentrated are qualitative ways to describe concentration
• dilute solution has relatively small concentration of solute • concentrated solution has high concentration of solute
• Quantitative expressions require information as masses, moles, or liters of the solute, solvent, or solution
• most commonly used expressions are: • mass percentage • mole fraction • molarity • molality
Mass Percentage, ppm, and ppb
• Mass percentage by definition: mass % of component = [(mass of component in soln/total mass of solution)]*100
• parts per million (ppm) expressed as number of mg of solute per kilogram of solution; by definition
• parts per million (ppm) of component = [(mass of component in soln/total mass of solution)]*106
• density of very dilute aqueous solution is similar to pure water (1g/ml)
• If density of solution is 1g/ml, then 1 ppm = 1 mg solute per liter of solution
• parts per billion (ppb) expressed as number of µg of solute per kilogram of solution
• By definition: • parts per billion (ppb) of component = [(mass of component in soln/total mass of solution)]*109
• If density of solution is 1g/ml, then 1 ppb = 1 µg solute per liter of solution
Mole Fraction, Molarity, and Molality
• Common expressions of concentration based on number of moles of one or more components
• Recall: mass can be converted to moles using the molar mass
• Recall: Mole fraction of component, X = moles of component/total moles of all components
• Note mole fraction has no units • Note mole fractions range from 0 to 1
• Recall: Molarity, M = moles of solute/litres of solution
• Note molarity will change with change in temperature (as solution volume increases or decreases)
• define molality (m), another concentration unit: molality, m = moles of solute/kilograms of solvent
• Molality not vary with temperature
• converting between molarity (M) & molality (m) requires density
• molarity & molality of dilute solutions are very similar
13.5 Colligative Properties
• Colligative properties depend number of solute particles
• 4 colligative properties: vapor pressure lowering (Raoult's Law), boiling point elevation, freezing point depression
• osmotic pressure, if time permits
Lowering the Vapor Pressure
• a volatile liquid in closed container
• after time equilibrium established between liquid & its vapor; partial pressure exerted by vapor is vapor pressure
• Nonvolatile solutes (no measurable vapor pressure) reduce ability of surface solvent molecules to escape liquid
• therefore, vapor pressure is lowered, amount of vapor pressure lowering depends on amount of solute
• Raoult’s law quantifies the extent to which nonvolatile solute lowers vapor pressure of solvent
• PA is vapor pressure with solute, P˚A is vapor pressure of pure solvent, &
is mole fraction of A
o
• ideal solution is one that obeys Raoult’s law: PA = XAP A
• Real solutions show ideal behavior when:
• solute concentration is low, solute & solvent similarly sized molecules, & similar types of inter-attractions
Boiling-Point Elevation
• nonvolatile solute lowers vapor pressure of a solution
• At normal boiling point of pure liquid, solution has vapor pressure less than 1 atm
• 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; ∆Tb = Kbm
• nature of solute (electrolyte vs. nonelectrolyte) will impact colligative molality of solute
Freezing-Point Depression
• When a solution freezes, crystals of almost pure solvent are formed first
• Solute molecules usually not soluble in solid phase of solvent
• the triple point occurs at lower temperature because of lower vapor pressure for solution
• melting-point (freezing-point) curve is vertical line from the triple point
• solution freezes at lower temperature (∆Tf) than pure solvent
• decrease in freezing point (∆Tf) is directly proportional to molality
• Kf is molal freezing-point-depression constant; ∆Tf = Kfm
• Values of Kf and Kb for most common solvents can be found in Table 13.4