Chapter 13 Solutions Intermolecular Forces & Solubility

Chapter 13 Solutions Intermolecular Forces & Solubility
Liquid Solutions Gas Solutions Solid Solutions The Solution Process Heats of Solution Heats of Hydration Solution Process & Change in Entropy 6/16/2018

Chapter 13 Solution as an Equilibrium Process Effect of Temperature
Effect of Pressure Concentration Molarity & Molality Parts of Solute by Parts of Solution Interconversion of Concentration Terms 6/16/2018

Chapter 13 Colligative Properties of Solutions
Nonvolatile Nonelectrolyte Solutions Solute Molar Mass Volatile Nonelectrolyte Solutions Strong Electrolyte Solutions Structure and Properties of Colloids 6/16/2018

Intermolecular Forces
Phases and Phase Changes (Chapter 12) are due to forces among the molecules Bonding (Intramolecular) forces and Intermolecuar forces arise from electrostatic attractions between opposite charges Bonding (Intramolecular) forces Attraction between Cations & Anions (ionic bonding) Electron pairs (covalent bonding) Metal Cations and Delocalized valence electrons (metallic bonding) Intermolecular forces Attractions between molecules as a result of partial charges Attractions between ions & molecules 6/16/2018

Bonding (Intramolecular) Relatively strong, larger charges close to each other Intermolecular Relatively weak involving smaller charges that are further apart Types (in decreasing order of bond energy) Ion-Dipole (strongest) Hydrogen bonding Dipole-Dipole Ion-Induced Dipole Dipole-Induced Dipole Dispersion (London) (weakest) 6/16/2018

Bonding (Intramolecular) Forces NonBonding (Intermolecular) Forces 6/16/2018

Intermolecular Forces & Solubility
Solutions – Homogeneous mixtures consisting of a solute dissolved in a solvent through the actions of intermolecular forces Solute dissolves in Solvent Distinction between Solute & Solvent not always clear Solubility – Maximum amount of solute that dissolves in a fixed quantity of solvent at a specified temperature “Dilute” & “Concentrated” are qualitative (relative) terms 6/16/2018

Substances with similar types of intermolecular forces (IMFs) dissolve in each other “Like Dissolves Like” Nonpolar substances (e.g., hydrocarbons) are dominated by dispersion force IMFs and are soluble in other nonpolar substances Polar substances have hydrogen bonding, ionic, and dipole driven IMFs (e.g., methanol) and would be soluble in polar substances (solvents) 6/16/2018

Ion-Dipole forces Principal forces involved in solubility of ionic compounds in water Each ion on crystal surface attracts oppositely charged end of water dipole These attractive forces overcome attractive forces between crystal ions leading to breakdown of crystal structure 6/16/2018

Hydrogen Bonding (N, O, F) Especially important in aqueous (water) solutions. Oxygen- and nitrogen-containing organic and biological compounds, such as alcohols, sugars, amines, amino acids. O & N are small and electronegative, so their bound H atoms are partially positive and can get close to negative O in the dipole water (H2O) molecule F H : O N 6/16/2018

Dipole-Dipole forces In the absence of H bonding, Dipole-Dipole forces account for the solubility of polar organic molecules, such as Aldehydes (R-CHO) in polar solvents such as chloroform (CHCl3) 6/16/2018

Ion-induced Dipole forces One of 2 types of charged-induced dipole forces (next slide – dipole-induced dipole) Rely on polarizability of components Ion’s charge distorts electron cloud of nearby nonpolar molecule Ion increases magnitude of any nearby dipole; thus contributes to solubility of salts in less polar solvents – LiCl in ethanol 6/16/2018

Dipole-induced dipole forces Intermolecular attraction between a polar molecule and the oppositely charged pole it induces in a nearby molecule Also based on polarizability, but are weaker than ion-induced dipole forces because the magnitude of charge is smaller – ions vs dipoles (coulombs law) Solubility of atmospheric gases (O2, N2, noble gases) have limited solubility in water because of these forces 6/16/2018

Dispersion forces (Instantaneous Dipoles) Contribute to solubility of all solutes in all solvents. Principal type of intermolecular force in solutions of nonpolar substances, such as petroleum and gasoline Keep cellular macromolecules in their biologically active shapes 6/16/2018

Liquid Solutions Solutions can be: Liquid Gaseous Solid
Physical state of Solvent determines physical state of solution Liquid solutions – forces created between solute and solvent comparable in strength (“like” dissolves “like”) Liquid – Liquid Alcohol/water – H bonds between OH & H2O Oil/Hexane – Dispersion forces Solid – Liquid Salts/Water – Ionic forces Gas – Liquid (O/H2O) – Weak intermolecular forces (low solubility, but biologically important) 6/16/2018

Liquid Solutions Gas Solutions
Gas – Gas Solutions – All gases are infinitely soluble in one another Gas – Solid Solutions – Gases dissolve into solids by occupying the spaces between the closely packed particles Utility – Hydrogen (small size) can be purified by passing an impure sample through a solid metal like palladium (forms Pd-H covalent bonds) Disadvantage – Conductivity of Copper is reduced by the presence of Oxygen in crystal structure (copper metal is transformed to Cu(I)2O 6/16/2018

Liquid Solutions Solid Solutions Solid – Solid Solutions
Alloys – A mixture with metallic properties that consists of solid phases of two or more pure elements, solid-solid solution, or distinct intermediate (heterogeneous) phases Alloys Types Substitutional alloys – Brass (copper & zinc), Sterling Silver(silver & copper), etc. substitute for some of main element atoms Interstitial alloys – atoms of another elements (usually a nonmetal) fill interstitial spaces between atoms of main element, e.g., carbon steel (C & Fe) 6/16/2018

Practice Problem A solution is ____________
a. A heterogeneous mixture of 2 or more substances b. A homogeneous mixture of 2 or more substances c. Any two liquids mixed together d. very unstable e. the answer to a complex problem Ans: b 6/16/2018

The Solution Process Elements of Solution Process
Solute particles must separate from each other Some solvent particles must separate to make room for solute particles Solute and Solvent particles must mix together Energy is absorbed when particles separate Energy is released when particles mix and attract each other Thus, the solution process is accompanied by changes in Enthalpy (sum of internal energy and product of pressure & volume ) and Entropy (number of ways the energy of system can be dispersed through motions of particles) 6/16/2018

Solution Process - Solvation
Solvation – Process of surrounding a solute particle with solvent particles Hydrated ions in solution are surrounded by solvent molecules in defined geometric shapes Orientation of solvent is different for cations and anions Cations attract the negative charge of the solvent Anions attract the positive charge of the solvent Hydration # for Na+ = 6 Hydration # for Cl- = 5 6/16/2018

Solution Process Solute v. Solvent
Coulombic - attraction of oppositely charged particles van der Waals forces: a. dispersion (London) b. dipole-dipole c. ion-dipole Hydrogen-bonding - attractive interaction of a Hydrogen atom and electronegative Oxygen, Nitrogen or Fluorine Hydration Energy attraction between solvent and solute Lattice Energy solute-solute forces 6/16/2018

Practice Problem Predict which solvent will dissolve more of the given solute: a. Sodium Chloride (solute) in Methanol or in 1-Propanol Ans: Methanol A solute tends to be more soluble in a solvent whose intermolecular forces are similar to those of the solute NaCl is an ionic solid that dissolves through ion-dipole forces Both Methanol & 1-Propanol contain a “Polar” -OH group The Hydrocarbon group in 1-Propanol is longer and can form only weak dispersive forces with the ions; thus it is less effective at substituting for the ionic attractions in the solute 6/16/2018

Practice Problem (con’t)
b. Ethylene Glycol (HOCH20CH20H) in Hexane (CH3(CH2)4CH3) or in Water (H2O) Ans: Very Soluble in Water Ethylene Glycol molecules have two polar –OH groups and the Molecules interact with each other through H-Bonding The Water H-bonds can substitute for the Ethylene Glycol H-Bonds better than they can with the weaker dispersive forces in Hexane (no H-Bonds) 6/16/2018

Practice Problem (con’t)
c. Diethyl Ether (CH3CH2OCH2CH3) in water or in Ethanol (CH3CH2OH) Ans: Ethanol Diethyl Ether molecules interact with each other through dipole-dipole and dispersive forces and can form H-Bonds to both H2O and Ethanol The ether is more soluble in Ethanol because Ethanol can form H-Bonds with the ether and its hydrocarbon group (CH3CH2) can substitute for the ether’s dispersion forces Water, on the other hand, can form H bonds with the ether, but it lacks any hydrocarbon portion, so it forms much weaker dispersion forces with the ether 6/16/2018

The Solution Process Solute particles absorb heat (endothermic) and separate from each other by overcoming intermolecular attractions Solvent particles absorb heat (endothermic) and separate from each other by overcoming intermolecular attractions Solute and Solvent particles mix (form a solution) by attraction and release energy - an exothermic process Total Enthalpy change is the “Heat of Solution” (∆Hsoln) Endothermic Endothermic Exothermic 6/16/2018

The Solution Process If (Hsolute + Hsolvent) < Hmix exothermic
Solution cycles and the enthalpy components of the heat of solution If (Hsolute + Hsolvent) < Hmix exothermic If (Hsolute + Hsolvent) > Hmix endothermic Exothermic solution process Solute Soluble Endothermic solution process Solute Insoluble Note: Recall Born-Haber Process – Chap 6 6/16/2018

Heats of Hydration Hsolvent & Hmix are difficult to measure individually Combined, these terms represent the Enthalpy change during “Solvation” – The process of surrounding a solute particle with solvent particles Solvation in “Water” is called “Hydration” Thus: Since: Then: Forming strong ion-dipole forces during hydration (Hmix <0) more than compensates for the braking of water bonds; thus the process of hydration is “exothermic”, i.e., 6/16/2018

Heat of Hydration & Lattice Energy
The energy required to separate an ionic solute (Hsolute) into gaseous ions is the lattice energy (Hlattice) This process requires energy input (endothermic) Thus, the heat of solution for ionic compounds in water combines the lattice energy (always positive) and the combined heats of hydration of cation and anion (always negative) Hsoln can be either positive or negative depending on the net values of Hlattice and Hhydr If the lattice energy is large relative to the hydration energy, the ions are likely to remain together and the compound will tend to be more insoluble 6/16/2018

Charge Density - Heats of Hydration
Heat of Hydration is related to “Charge Density” of ion Charge Density – ratio of ion’s charge to its volume Coulombs Law – the greater the ion’s charge and the closer the ion can approach the oppositely charged end of the water molecule, the stronger the attraction and, thus, the greater the charge density The higher the charge density the more negative Hhydr,, i.e., the more soluble Periodic Table Going down Group – same charge, increasing size Charge density & Heat of Hydration decrease Going across Period – greater charge, smaller size Charge density & Heat of Hydration increase 6/16/2018

Heat of Hydration & Lattice Energy
Hydration vs. Lattice Energies The solubility of an ionic solid depends on the relative contribution of both the energy of hydration (Hhydr) and the lattice energy (Hlattice) Lattice Energy Small size + high charge = high lattice E For a given charge, Lattice energy (Hlattice) decreases as the radius of ions increases, increasing solubility Energy of Hydration Energy of Hydration (Hhydr) also depends on ionic radius Small ions, such as Na1+ or Mg2+ have concentrated electric charge and a strong electric field that attracts water molecules High Charge Density (more negative Hhydr) higher solubility 6/16/2018

Heats of Hydration Cation Radius (pm) Hhydr (kJ/Mol) Anion Hhydr Na+
102 -410 F- 119 -431 K+ 138 -336 OH- -460 Rb+ 152 -315 NO31- 165 -314 Cs+ 167 -282 Cl- -313 Mg2+ 72 -1903 CN1- 177 -342 Ca2+ 100 -1591 N31- 181 -298 Sr2+ 118 -1424 Br- 182 -284 Ba2+ 135 -1317 SH1- 193 I- 206 -247 BF41- 215 -223 ClO41- 226 -235 CO32- 164 -1314 S2- 170 -1372 SO42- 244 -1059 6/16/2018

Solubility - Hhydr vs Hlattice
Relative Solubilities – Alkaline Earth Hydroxides Mg(OH)2 < Ca(OH)2 < Sr(OH)2 < Ba(OH)2 Lattice energy decreases as the radius of alkaline earth ion increases from Mg2+ to Ba2+  Lattice energy alone would suggest solubility to increase down the group from Magnesium Hydroxide to Barium Hydroxide Energy of Hydration is greatest for small ions (Mg2+)  Energy of Hydration alone would suggest solubility to decrease down the group from Mg(OH)2 to Ba(OH)2 6/16/2018

Relative Solubilities - Alkaline Earth Sulfates MgSO4 > CaSO4 > SrSO4 > BaSO4 Lattice energy depends on the sum of the anion/cation radii Since the Sulfate (SO42-) ion is much larger than the hydroxide ion (244 pm vs 119 pm), the percent change in lattice energy going from Magnesium to Barium in the sulfates is smaller than for the hydroxides The energy of hydration for the cations decreases by a greater amount going from Mg to Ba in the sulfates relative to the hydroxides 6/16/2018

Energy Solubility Hydration Energy Solubility Mg++ Alkaline Earth Hydroxides Alkaline Earth Sulfates (Hlattice) (HHydration) High Low High High Low High Ionic Size Ba++ Low High Low Low High Low Hydroxides Down group, the lattice energy of alkaline earth hydroxides decreases more rapidly than does the energy of hydration  Lattice-energy factor (Hlattice) dominates Ba(OH)2 more soluble than Mg(OH)2 Sulfates The Sulfate (SO42-) ion is much larger than the hydroxide ion (244 pm vs 119 pm) The percent change in lattice energy going from Magnesium to Barium in the sulfates is smaller than for the hydroxides The energy of hydration for the cations decreases by a greater amount going from Mg to Ba in the sulfates relative to the hydroxides (Hhyd) dominates MgSO4 is more soluble than BaSO4 6/16/2018

Solubility – Solution Temperatures
Solubilities and Solution temperatures for 3 salts NaCl NaOH NH4NO3 The interplay between (Hlattice) and (Hhydr) leads to different values of (Hsoln) The resulting temperature of the solution is a function of the relative value of (Hsoln) Neg (Hot) Pos (Cold) Hlattice (Hsolute) is always positive for ionic compounds Combined ionic heats of hydration (Hhydr ) are always negative Hot NaOH Hot Hsoln = kJ/mol NaCl Neutral Hsoln = 3.9 kJ/mol NH4NO3 Cold Hsoln = 25.7 kJ/mol 6/16/2018

Predicting Solubility
Compatible vs Incompatible Intermolecular Forces (IMFs) C6H14 + C10H22 compatible; both dispersion forces & non-polar molecule (soluble) C6H14 + H2O incompatible; dispersion vs polar and H-bonding (insoluble) H2O + CH3OH compatible; both H-bonding (soluble) NaCl + H2O compatible; ion-dipole (soluble) NaCl + C6H6 incompatible; ion vs dispersion (insoluble) CH3Cl + CH3COOH compatible; both polar covalent & dipole-dipole (soluble) 6/16/2018

Practice Problem Predicting Solubility Which is more soluble in water?
NaCl vs C6H6 CH3OH vs C6H6 CH3OH vs NaI CH3OH vs CH3CH2CH2OH BaSO4 vs MgSO4 BaF vs MgF Ans: NaCl Ans: CH3OH Ans: NaI Ans: CH3OH Ans: Mg(SO4) Ans: BaF 6/16/2018

Writing Solution Equations
1. Write out chemical equation for the dissolution of NaCl(s) in water. NaCl(s)  Na+(aq) + Cl-(aq) 2. Write out chemical equation for the dissolution of tylenol (C8H9NO2)in water C8H9NO2(s)  C8H9NO2(aq) H2O H2O 6/16/2018

Solution Process - Entropy
The heat of solution is one of two factors that determine whether a solute dissolves in a solvent The other factor concerns the natural tendency of a system to spread out, distribute, or disperse its energy in as many ways as possible Entropy (S) is a Thermodynamic variable directly related to the number of ways a system can distribute its energy Entropy will be discussed in more detail in Chapter 18 Entropy is related to freedom of motion of the particles in the system and the number of ways they can be arranged 6/16/2018

Gas vs Liquid vs Solid A liquid has a higher Entropy (S) than a solid because the particles have higher freedom of motion Sliquid > Ssolid Similarly, a gas has higher entropy than a liquid Sgas > Sliquid The change in entropy when a liquid changes to a gas is positive Svap > 0 A solution usually has a higher entropy than a pure solute and a pure solvent. The solution provides more ways to distribute energy Particles in a solution have a higher freedom of motion with increased possibilities for interactions between molecules 6/16/2018

Sodium Chloride Solution in Hexane vs Toluene in Hexane The ion-dipole forces in NaCl are weak and incompatible with the dispersion forces in Hexane (Hydrocarbon – C6H14) Hmix (exothermic) is much smaller than Hsolute (endothermic) Hsoln >> 0  NaCl insoluble Solution does not form because the entropy increase that would accompany the mixing of solute & solvent is much smaller than the enthalpy increase required to separate The dispersive forces in hexane and toluene are compatible Hmix dominates Hsolute & Hsolvent  soluble 6/16/2018

Solubility as Equilibrium Process
Ions disaggregate and become dispersed in solvent Some undissolved solids collide with undissolved solute and recrystallize Equilibrium is reached when rate of dissolution equals rate or recrystallization Saturation – Solution contains the maximum amount of dissolved solute at a given temperature Unsaturation – Solution contains less than maximum amount of solute – more solute could be dissolved! Supersaturation – Solution contains more dissolved solute than equilibrium amount Supersaturation can be produced by slowly cooling a heated equilibrium solution 6/16/2018

Comparison of Unsaturated and Saturated Solutions
6/16/2018

Solubility Equilibrium
A(s)  A(aq) A(s) = solid lattice form A(aq) = solvated solution form Rate of dissolution = rate of crystallization Amount of solid remains constant Dynamic Equilibrium – As many particles going into solution as coming out of solution 6/16/2018

Gas Solubility – Temperature
Gas particles (solute) are already separated; thus, little heat required for disaggregation Hsolute  0 Heat of hydration of a gaseous species is always exothermic (Hhydration < 0) Hsoln = Hsolute + Hhydration < 0 Solute(g) + solvent(l)  sat’d soln + heat Gas molecules have weak intermolecular forces and the intermolecular forces between a gas and solvent are also weak As temperature rises, the kinetic energy of the gas molecules also increases allowing them to escape, lowering concentration, i.e. Solubility of a gas decreases with increasing temperature 6/16/2018

Salt Solubility – Temperature
Most solids are more soluble at higher temperatures Most ionic solids have a positive Hsoln because (Hlattice) is greater than (HHydr) Note exception with Ce2(SO4)3 Hsoln is positive, heat is absorbed to form solution If heat is added, the rate of solution should increase 6/16/2018

Effect of Pressure on Gas Solubility
At a given pressure, the same number of gas molecules enter and leave a solution per unit time – equilibrium As partial pressure (P) of gas above solution increases, the concentration of the molecules of gas increases in solution 6/16/2018

Effect of Pressure on Solutions
Pressure changes do not effect solubility of liquids and solids significantly aside from extreme P changes (vacuum or very high pressure) Gases are compressible fluids and therefore solubilities in liquids are Pressure dependent Solubility of a gas in a liquid is proportionaL () to the partial pressure of the gas above the solution Quantitatively defined in terms of Henry’s Law S = kHP S = solubility of gas (mol/L) Pg = partial pressure of gas (atm) KH = Henry’s Law Constant (mol/L-atm) 6/16/2018

Practice Problem The solubility of gas A (Mm = 80 g/mol) in water is 2.79 g/L at a partial pressure of 0.24 atm. What is the Henry’s law constant for A (atm/M)? 6/16/2018

Practice Problem If the solubility of gas Z in water is 0.32 g/L at a pressure of atm, what is the solubility (g/L) of Z in water at a pressure of atm? 6/16/2018

moles of solute + moles of solvent
Concentration Definitions Concentration Term Ratio amount (mol) of solute volume (L) of solution Molarity (M) amount (mol) of solute mass (kg) of solvent Molality (m) mass of solute mass of solution Parts by mass volume of solute volume of solution Parts by volume amount (mol) of solute moles of solute + moles of solvent Mole Fraction (X) 6/16/2018

Practice Problem The molality of a solution is defined as _______ a. moles of solute per liter of solution. b. grams of solute per liter of solution. c. moles of solute per kilogram of solution. d. moles of solute per kilogram of solvent. e. the gram molecular weight of solute per kilogram of solvent. Ans: d 6/16/2018

Practice Problem Fructose, C6H12O6 (FW = g/mol, is a sugar occurring in honey and fruits. The sweetest sugar, it is nearly twice as sweet as sucrose (cane or beet sugar). How much water should be added to 1.75 g of fructose to give a m solution? 6/16/2018

Colligative Properties of Solutions
Presence of Solute particles in a solution changes the physical properties of the solution The number of particles dissolved in a solvent also makes a difference in four (4) specific properties of the solution known as – “Colligative Properties” Vapor Pressure Boiling Point Elevation Freezing Point Depression Osmotic Pressure 6/16/2018

Phase Diagram for pure solvent and for Solution Pure Liquid Solvent Triple Point Of Solvent Solution Pure Solid Solvent Triple Point Of Solution 6/16/2018

Colligative properties deal with the nature of a solute in aqueous solution and the extent of the dissociation into ions Electrolyte – Solute dissociates into ions and solution is capable of conducting an electric current Strong Electrolyte – Soluble salts, strong acids, and strong bases dissociate completely; thus, the solution is a good conductor Weak Electrolyte – polar covalent compounds, weak acids, weak bases dissociate weakly and are poor conductors Nonelectrolyte – Compounds that do not dissociate at all into ions (sugar, alcohol, hydrocarbons, etc.) are nonconductors 6/16/2018

Prediction of the magnitude of a colligative property Solute Formula Each mole of a nonelectrolyte yields 1 mole of particles in the solution Ex M glucose contains 0.35 moles of solute particles (glucose molecules) per liter Each mole of strong electrolyte dissociates into the number of moles of ions in the formula unit Ex. 0.4 M Na2SO4 contains 0.8 mol of Na+ ions and 0.4 mol of SO42- ions (total 1.2 mol of particles) per liter of solution. 6/16/2018

Vapor Pressure Vapor pressure (Equilibrium Vapor Pressure):
The pressure exerted by a vapor at equilibrium with its liquid in a closed system Vapor Pressure increases with increasing temperature nonvolatile nonelectrolyte (ex. sugar) & pure solvent The vapor pressure of a solution of a nonvolatile nonelectrolyte (solute) is always lower than the vapor pressure of the pure solvent Presence of solute particles reduces the number of solvent vapor particles at surface that can vaporize At equilibrium, the number of solvent particles leaving solution are fewer than for pure solvent, thus, the vapor pressure is less Entropy of solution is always greater than for solvent; thus, it has less tendency to vaporize to gain entropy 6/16/2018

Vapor Pressure The vapor pressure of the solvent above the solution (Psolvent) equals the mole fraction of solvent in the solution (Xsolvent) times the vapor pressure of the pure solvent (Raoult’s Law) Psolvent = Xsolvent x Posolvent In a solution, the mole fraction of the solvent (Xsolvent) is always less than 1; thus the partial pressure of the solvent above the solution (Psolvent) is always less the partial pressure of the pure solvent Posolvent Ideal Solution – An ideal solution would follow Raoult’s law for any solution concentration Most gases in solution deviate from ideality Dilute solutions give good approximation of Raoult’s law 6/16/2018

Vapor Pressure A solution consists of Solute & Solvent
The sum of their mole fractions equals 1 The magnitude of P (vapor pressure lowering) equals the mole fraction of the solute times the vapor pressure of the pure solvent 6/16/2018

Vapor Pressure Vapor Pressure & Volatile Nonelectrolyte Solutions
The vapor now contains particles of both a volatile solute and the volatile solvent From Dalton’s Law of Partial Pressures The presence of each volatile component lowers the vapor pressure of the other by making each mole fraction less than 1 6/16/2018

Mole fractions in vapor are different
Example Problem Equi-molar solution of Benzene and Toluene, both nonelectrolyte hydrocarbons Benzene lowers the vapor pressure of Toluene and Toluene lowers the vapor pressure of Benzene Vapor composition vs Solution composition Mole fractions in vapor are different 6/16/2018

Boiling Point Elevation
Boiling Point (Tb) of a liquid is the temperature at which its vapor pressure equals the external (atmospheric usually) pressure The vapor pressure of a solution (solvent + solute) is lower than that of the pure solvent at any temperature Thus, the difference between the external (atmospheric) pressure and the vapor pressure of the solution is greater than the difference between external pressure and the solvent vapor pressure - Psoln > Psolvent The boiling point of the solution will be higher than the solvent because additional energy must be added to the solution to raise the vapor pressure of the solvent (now lowered) to the point where it again matches the external pressure The boiling point of a concentrated solution is greater than the boiling point of a dilute solution 6/16/2018

Boiling Point Elevation
The magnitude of the boiling point elevation is proportional to the Molal concentration of the solute particles Note the use of “Molality” Molality is related to mole fraction, thus particles of solute Molality also involves “Mass of Solvent”; thus, not affected by temperature 6/16/2018

Freezing Point Depression
Recall: The vapor pressure of a solution is lower than the vapor pressure of the pure solvent The lower vapor pressure of a solution is the result of fewer solvent molecules being able to leave the liquid phase because of the competition with solute molecules for space at the surface In solutions with nonvolatile solutes, only solvent molecules can vaporize; thus, only solvent molecules can solidify (freeze) The formation of solid solvent particles leaves the solution more concentrated (less solvent available) At the freezing point of a solution, there is an equilibrium established between solid solvent and liquid solvent There is also an equilibrium between the number of particles of solvent leaving the solution and the number of particles returning to the solution 6/16/2018

The vapor pressure of a pure solvent at its freezing point, i.e., an equilibrium solution of solid and liquid solvent, is higher than the vapor pressure of a solution of a solute in this solvent The vapor pressure of a solvent in a solution at its freezing point has been lowered from its value as a pure solvent until it equals that of the solution Since vapor pressure is a function of temperature, the solution, with its lower vapor pressure, will freeze at a lower temperature than the pure solvent 6/16/2018

The Freezing Point depression has the magnitude proportional to the Molal concentration of the solute 6/16/2018

The van’t Hoff Factor Colligative properties depend on the relative number of solute to solvent “particles” In strong electrolytes solutions, the solute formula specifies the number of particles affecting the colligative property Ex. The BP elevation of a 0.5 m NaCl soln would be twice that of a 0.5 m glucose soln because NaCL dissociates into “2” particles per formula unit, where glucose produces “1” particle per formula unit The van’t Hoff factor (i) is the ratio of the measured value of the colligative property, e.g. BP elevation, in the electrolyte solution to the expected value for a nonelectrolyte solution 6/16/2018

The van’t Hoff Factor To calculate the colligative properties of strong electrolyte solutions, incorporate the van’t Hoff factor into the equation CH3OH(l)  CH3OH(aq) 1:1, i = 1 NaCl(s)  Na+(aq) + Cl-(aq) 2:1, i = 2 CaCl2(s)  Ca2+(aq) Cl-(aq) 3:1, i = 3 Ca3(PO4)2(s)  3 Ca2+(aq) PO43-(aq) 5:1, i = 5 Ex. Freezing Point Depression with van’t Hoff factor for Ca3(PO4)2(s) 6/16/2018

Colligative Mathematical Property Relation 1. vapor pressure (Raoult’s Law) 2. freezing pt depression 3. boiling pt elevation 4. osmotic pressure 6/16/2018

Practice Problem What is the boiling point of m aqueous calcium chloride, CaCl2? 6/16/2018

Practice Problem What is the freezing point of a 0.25 m solution of glucose in water (Kf for water is 1.86°C/m.)? a. 0.93°C b. –0.93°C c. 0.46°C d. –0.46°C e. 0.23°C Ans: d 6/16/2018

Practice Problem What is the freezing point of g of glycerol (C3H8O3) in g of water? 6/16/2018

Practice Problem What is the molar mass (Mm) of Butylated hydroxytoluene (BHT) if a solution of g of BHT in g of benzene (Kf = oC/m; Tf = oC) had a freezing point of oC? 6/16/2018

Colloids Suspensions vs Mixtures vs Colloids
A Heterogeneous mixture – fine sand suspended in water – consists of particles large enough to be seen by the naked eye, clearly distinct from surrounding fluid A Homogeneous mixture – sugar in water – forms a solution consisting of molecules distributed throughout and indistinguishable from the surrounding fluid Between these extremes is a large group of mixtures called colloidal dispersions – Colloids Colloid particles are larger than simple molecules, but small enough to remain in suspension and not settle out Colloids range in size from 1 – 1000 nm 6/16/2018

Colloids Colloids have tremendous surface areas, which allows many more interactions to exert an large total adhesive force, which attracts other particles Many useful applications Colloids are classified according to whether the dispersed and dispersing substances are gases, liquids, or solids 6/16/2018

Osmosis Osmosis, a colligative property, is the movement of solvent particles through a semipermeable membrane separating a solution of one concentration from a solution of another concentration, while the solute particles stay within their respective solutions Many organisms have semipermeable membranes that regulate internal concentrations. Water can be purified by “reverse osmosis” The direction of flow of the solvent is from the solution of lower concentration to the solution of higher concentration increasing its volume, decreasing concentation Osmotic pressure is defined as the amount of pressure that must be applied to the higher concentration solution to prevent the dilution and change in volume 6/16/2018

Osmosis 6/16/2018

Osmosis (2-Compartment Chamber)
1 2                            Semipermeable membrane = solute molecule  = solvent molecule Osmosis = solvent flow across membrane Net movement of solvent is from 2 to 1; pressure increases in 1 PA,1 < PA,2, osmosis follows solvent vapor pressure Osmotic Pressure = pressure applied to just stop osmosis 6/16/2018

Molarity (M) (moles per liter of solution)
Osmotic Pressure At equilibrium the rate of solvent movement into the more concentrated solution is balanced by the rate at which the solvent returns to the less concentrated solution. The pressure difference at equilibrium is the Osmotic Pressure (), which is defined as the applied pressure required to prevent the net movement of solvent from the less concentrated solution to the more concentrated solution Osmotic Pressure () is proportional to the number of solute particles in a given volume of solution, i.e., Molarity (M) (moles per liter of solution) 6/16/2018

Osmotic Pressure Osmotic Pressure / Solution Properties
The volume of the concentrated solution will increase as the solvent molecules pass through the membrane into the solution decreasing the concentration This change in volume is represented as a change in height of the liquid in a tube open to the externally applied pressure This height change (h), the solution density (d), and the acceleration of gravity (g) can also be used to compute the Osmotic Pressure ()  = gdh  = g(m/s2)d(kg/m3)h(m) = pascals(kg/ms2) 1 pascal = x 10-3 torr(mm)= x 10-6 atm 6/16/2018

Reverse Osmosis Membranes with pore sizes of 0.1 – 10 m are used to filter colloids and microorganisms out of drinking water In reverse osmosis, membranes with pore sizes of – 0.01 m are used to remove dissolved ions A pressure greater than the osmotic pressure is applied to the concentrated solution, forcing the water back through the membrane, filtering out the ions Removing heavy metals from domestic water supplies and desalination of sea water are common uses of reverse osmosis 6/16/2018

Practice Problem How much potable water can be formed per 1,000 L of seawater (~0.55 M NaCl) in a reverse osmosis plant that operates with a hydrostatic pressure of 75 atm at 25 oC? 6/16/2018

Practice Problem What is the osmotic pressure (mm Hg) associated with a M aqueous calcium chloride, CaCl2, solution at 25 oC? a mm Hg b mm Hg c mm Hg d mm Hg e mm Hg 6/16/2018

Practice Problem Consider the following dilute NaCl(aq) solutions
a. Which one will boil at a higher temperature? b. Which one will freeze at a lower temperature? c. If the solutions were separated by a semi permeable membrane that allowed only water to pass, which solution would you expect to show an increase in the concentration of NaCl? Ans: a. B (increase temperature (VP) to match atmospheric pressure b. B (decrease temperature to reduce vapor pressure of solvent) c. A (The movement of solvent between solutions is from the solution of lower concentration to the solution of higher concentration) 6/16/2018

Practice Problem Consider the following three beakers that contain water and a non-volatile solute. The solute is represented by the orange spheres. a. Which solution would have the highest vapor pressure? b. Which solution would have the lowest boiling point? c. What could you do in the laboratory to make each solution have the same freezing point? a. A A Condense “A” by ½ 6/16/2018

Practice Problem Caffeine, C8H10N4O2 (FW = g/mol), is a stimulant found in tea and coffee. A Sample of the substance was dissolved in 45.0 g of chloroform, CHCl3, to give a m solution. How many grams of caffeine were in the sample? Ans: 6/16/2018

Practice Problem A solution contains g of a molecular compound in g of ethanol. The molality of the solution is m. Calculate the molecular weight of the compound 6/16/2018

Practice Problem What is the vapor pressure (mm Hg) of a solution of g of urea [(NH2)2CO, FW = 60.0 g/mol] in 3.00 g of water at 25 oC? What is the vapor pressure lowering of the solution? The vapor pressure of water at 25 oC is 23.8 mm Hg 6/16/2018

Practice Problem What is the vapor pressure lowering in a solution formed by adding 25.7 g of NaCl (FW = g/mol) to 100. g of water at 25 oC? (The vapor pressure of pure water at 25 oC is 23.8 mm Hg) 6/16/2018

Practice Problem A g sample of an ionic compound with the formula Cr(NH3)5Cl3 (FW = g/mol) was dissolved in water to give 25.0 mL of solution at 25 oC. The osmotic pressure was determined to be 119 mm Hg. How many ions are obtained from each formula unit when the compound is dissolved in water? 6/16/2018

Equation Summary Component Enthalpies of Heat of Solution
Component Enthalpies of Ionic Compound Heat of Soln Relating Gas Solubility to its Partial Pressure (Henry’s Law) Relationship between vapor pressure & mole fraction 6/16/2018

Equation Summary Vapor Pressure Lowering Boiling Point Elevation
Freezing Point Depression Osmotic Pressure 6/16/2018

amount (mol) of solute + amount (mol) of solvent
Equation Summary Ratio amount (mol) of solute volume (L) of solution Concentration Term Molarity (M) amount (mol) of solute mass (kg) of solvent Molality (m) mass of solute mass of solution Parts by mass volume of solute volume of solution Parts by volume Mole Fraction (X) amount (mol) of solute amount (mol) of solute + amount (mol) of solvent 6/16/2018

Chapter 13 Solutions Intermolecular Forces & Solubility

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Chapter 13 Solutions Intermolecular Forces & Solubility

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