Interactions and Solubility: Intermolecular Forces and Solutions

1. 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

2. 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

3. Chapter 13 • Colligative Properties of Solutions • Nonvolatile Nonelectrolyte Solutions • Solute Molar Mass • Volatile Nonelectrolyte Solutions • Strong Electrolyte Solutions • Structure and Properties of Colloids

4. Intermolecular Forces • Phases and Phase Changes (Chapter 12) are due to forces among the molecules • Bonding (Intramolecular) forces and Intermolecuarforces 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

5. Intermolecular Forces • 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. Intermolecular Forces Bonding (Intramolecular) Forces NonBonding (Intermolecular) Forces

7. Intermolecular Forces & Solubility • Solutions – Homogeneous mixtures consisting of a solutedissolved in a solventthrough the actions of intermolecularforces • 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

8. Intermolecular Forces & Solubility • 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)

9. Intermolecular Forces & Solubility • Intermolecular Forces • 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

10. Intermolecular Forces & Solubility • Intermolecular Forces • 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 : : : H N H O H F

11. Intermolecular Forces & Solubility • Intermolecular Forces • 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)

12. Intermolecular Forces & Solubility • Intermolecular Forces • 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

13. Intermolecular Forces & Solubility • Intermolecular Forces • 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

14. Intermolecular Forces & Solubility • Intermolecular Forces • 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

15. 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)

16. 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

17. 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)

18. 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

19. 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 absorbedwhen particles separate • Energy is releasedwhen 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)

20. 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

21. 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

22. 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

23. 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)

24. Practice Problem (con’t) c. Diethyl Ether (CH3CH2OCH2CH3) in water or inEthanol (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

25. The Solution Process • Solute particles absorb heat (endothermic) and separate from each other by overcoming intermolecular attractions • Solventparticles 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 exothermicprocess • Total Enthalpy change is the “Heat of Solution” (∆Hsoln) Endothermic Endothermic Exothermic

26. The Solution Process • Solution cycles and the enthalpy components of the heat of solution If(Hsolute + Hsolvent) < Hmixexothermic If(Hsolute + Hsolvent) > Hmixendothermic Exothermic solution process Solute Soluble Endothermic solution process Solute Insoluble Note: Recall Born-Haber Process – Chap 6

27. Heats of Hydration • Hsolvent& Hmixare difficult to measure individually • Combined, these terms represent the Enthalpy change during “Solvation” – The process of surrounding a solute particle with solvent particles • Solvationin “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.,

28. 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) • Hsolncan be either positiveor negativedepending 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

29. 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

30. 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

31. Heats of Hydration

32. 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

33. Solubility - Hhydr vs Hlattice • 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 smallerthan for the hydroxides • The energy of hydration for the cations decreasesby a greateramount going from Mg to Ba in the sulfates relative to the hydroxides

34. Solubility - Hhydr vs Hlattice Solubility Lattice Energy Solubility Hydration Energy Solubility Mg++ Alkaline Earth Sulfates Alkaline Earth Hydroxides (Hlattice) (HHydration) High Low High High Low High Ionic Size High Low Ba++ High Low Low Low • Hydroxides • Downgroup, the lattice energyof alkaline earth hydroxides decreasesmore rapidlythan 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 largerthan the hydroxide ion (244 pm vs 119 pm) • The percent change in lattice energygoing from Magnesium to Barium in the sulfates is smaller than for the hydroxides • The energy of hydration for the cationsdecreasesby a greateramount going from Mg to Ba in the sulfates relative to the hydroxides (Hhyd) dominates • MgSO4 is more soluble than BaSO4

35. 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 = -44.5 kJ/mol NH4NO3 Cold NaCl Neutral Hsoln = 25.7 kJ/mol Hsoln = 3.9 kJ/mol

36. 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)

37. Practice Problem Predicting Solubility Which is more soluble in water? • NaClvs C6H6 • CH3OH vs C6H6 • CH3OH vsNaI • CH3OH vs CH3CH2CH2OH • BaSO4vs MgSO4 • BaFvsMgF Ans: NaCl Ans: CH3OH Ans: NaI Ans: CH3OH Ans: Mg(SO4) Ans: BaF

38. 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

39. 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 Thermodynamicvariable 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 tofreedom of motionof the particles in the system and the number of ways they can be arranged

40. Solution Process - Entropy • 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

41. Solution Process - Entropy • 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 • HmixdominatesHsolute & Hsolvent soluble

42. Solubility as Equilibrium Process • Ions disaggregate and become dispersed in solvent • Some undissolved solids collide with undissolved solute and recrystallize • Equilibriumis 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

43. Comparison of Unsaturated andSaturated Solutions

44. Solubility Equilibrium • 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

45. 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 alwaysexothermic(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

46. Salt Solubility – Temperature • Most solids are more soluble at higher temperatures • Most ionic solids have a positiveHsolnbecause(Hlattice)is greater than(HHydr) • Note exception with Ce2(SO4)3 • Hsolnis positive,heat is absorbedto form solution • If heat is added, the rate of solution should increase

47. 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, theconcentrationof the molecules of gas increases in solution

48. Effect of Pressure on Solutions • Pressure changes do not effect solubility ofliquidsand solidssignificantly 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)

49. 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)?

50. Practice Problem If the solubility of gas Z in water is 0.32 g/L at a pressure of 1.0 atm, what is the solubility (g/L) of Z in water at a pressure of 0.0063 atm?