CH1810 Lecture #1 Solutions of Ionic Compounds

Transcription

1 CH1810 Lecture #1 Solutions of Ionic Compounds

2 Solutions Homogeneous mixtures are called solutions. The component of the solution that changes state is called the solute. The component that keeps its state is called the solvent. If both components start in the same state, the major component is the solvent.

3 Kinds of Solutions Gas in gas Gas in liquid Gas in solid Liquid in liquid Liquid in solid Solid in liquid Solid in solid Air (O2 in N2) Carbonated water (CO2 in H2O) H2 in palladium metal Gasoline, tequila Dental amalgam (Hg in Ag) Salt water (NaCl in H2O) Sterling silver (Cu in Ag)

4 What Happens When a Solute Dissolves? There are attractive forces between the solute particles holding them together. There are also attractive forces between the solvent molecules. When we mix the solute with the solvent, there are attractive forces between the solute particles and the solvent molecules. If the attractions between solute and solvent are strong enough, the solute will dissolve.

5 Attractive Forces and Solubility

6 Hierarchy of Intermolecular Forces Charged particles Molecules containing O-H, N-H, or F-H Bonds H-bonding Polar Molecules Dipole forces All Molecules Dispersion forces

7 Attractive Forces and Solubility Solubility depends, in part, on attractive forces between solute and solvent molecules Like dissolves like. Polar substances dissolve in polar solvents hydrophilic groups = OH, CHO, C=O, COOH, NH 2, Cl Nonpolar molecules dissolve in nonpolar solvents hydrophobic groups = C-H, C-C

8 Immiscible Liquids Pentane, C 5 H 12, is a nonpolar molecule. Water is a H 2 O, is a polar molecule. The attractive forces between the water molecules is much stronger than their attractions for the pentane molecules. The result is the liquids are immiscible* C5H12 H2O *Miscible liquids will dissolve in each other.

9 Nonpolar Solvents

10 Polar Solvents Dichloromethane (methylene chloride) Ethanol (ethyl alcohol) Water

11 Practice Choose the substance in each pair that is more soluble in water a) CH3OH CH3CHF2 b) CH3CH2CH2CH2CH3 CH3Cl

12 Salt vs. Sugar Dissolved in Water Ionic compounds dissociate into ions when they dissolve. Molecular compounds do not dissociate into ions when they dissolve.

13 Dissociation vs Ionization When ionic compounds dissolve in water, the anions and cations are separated from each other. This is called dissociation. Na 2 S(aq) 2 Na + (aq) + S 2- (aq) When compounds containing containing polyatomic ions polyatomi dissociate, the polyatomic group stays together as one ion. Na 2 SO 4 (aq) 2 Na + (aq) + SO 4 2 (aq) en strong When strong acids dissolve in water, in the water, molecule the ionizes into H+ and anions. H 2 SO 4 (aq) 2 H + (aq) + SO 4 2 (aq)

14 Energy Changes During Dissolution of Ionic Solids

15 Dissolution of Ionic Compounds

16 Temperature and Solubility of Ionic Compounds

17 Lewis Theory Predictions for Ionic Bonding Lewis theory predicts the number of electrons a metal atom should lose or a nonmetal atom should gain. This allows us to predict the formulas of ionic compounds that result. It also allows us to predict the relative strengths of the resulting ionic bonds from Coulomb s Law.

18 Energetics of Ionic Bond Formation The ionization energy of a metal is endothermic: Na(s) Na + (g) + 1 e ΔH = +496 kj/mol The electron affinity of a nonmetal is exothermic: ½Cl2(g) + 1 e Cl (g) ΔH = 244 kj/mol Therefore the formation of the ionic compound should be endothermic. But the heat of formation of most ionic compounds is exothermic and generally large. Na(s) + ½Cl2(g) NaCl(s) ΔH f = 411 kj/mol Why?

19 Ionic Bonding & the Crystal Lattice The extra energy that is released comes from the formation of a structure in which every cation is surrounded by anions. This structure is called a crystal lattice. The crystal lattice is held together by electrostatic attractions. The crystal lattice maximizes these attractions between cations and anions, leading to the most stable arrangement.

20 Lattice Energy

21 Lattice Energy The extra stability that accompanies the formation of the crystal lattice is measured as the lattice energy. The lattice energy is the energy released when the solid crystal forms from separate ions in the gas state 1) always exothermic 2) Can be calculated from knowledge of other processes Lattice energy depends directly on size of charges and inversely on distance between ions. U = k(q 1)(Q2) d

22 Lattice Energy vs. Ion Size

23 Born-Haber Cycle

24 ionization electron affinity bond breaking lattice energy sublimation Born-Haber Cycle

25 Born-Haber Cycle 1) Na(s) + ½Cl 2 (g) Na(g) + ½Cl 2 (g) (sublimation) + H sub ) Na(g) + ½Cl 2 (g) Na(g) + Cl (g) (bond energy) + ½H BE 3) Na(g) + Cl (g) Na + (g) + Cl (g) (ionization) +H IE1 4) Na + (g) + Cl (g) Na + (g) + Cl - (g) (electron affinity) H EA (lattice energy) 5) Na + (g) + Cl (g) NaCl(s) - U = ΔHf U = H f ½ H BE H EA H sub H IE1

26 Trends in Lattice Energy Ion Charge The force of attraction between oppositely charged particles is directly proportional to the product of the charges. Larger charge means the ions are more strongly attracted. larger charge stronger attraction stronger attraction larger lattice energy Of the two factors, ion charge is generally more important Lattice Energy = 910 kj/mol Lattice Energy = 3414 kj/mol

27 Trends in Lattice Energy Ion Magnitude The force of attraction between charged particles is inversely proportional to the distance between them. Larger ions mean the center of positive charge (nucleus of the cation) is farther away from the negative charge (electrons of the anion). larger ion weaker attraction weaker attraction smaller lattice energy

28 Lattice Energies of Some Ionic Solids (kj/mole) Anions Cations F- Cl- Br- I- O 2- Li ,925 Na ,695 K ,360 Be 2+ 3,505 3,020 2,914 2,800 4,443 Mg 2+ 2,957 2,524 2,440 2,327 3,791 Ca 2+ 2,630 2,258 2,176 2,074 3,401 Al 3+ 5,215 5,492 5,361 5,218 15,916

29 Order the following ionic compounds in order of increasing magnitude of lattice energy: CaO, KBr, KCl, SrO First examine the ion charges and order by sum of the charges (KBr, KCl) < (CaO, SrO) Then examine the ion sizes of each group and order by radius larger < smaller KBr < KCl < SrO < CaO

30 Order the following ionic compounds in order of increasing magnitude of lattice energy: MgS, NaBr, LiBr, SrS First examine the ion charges and order by sum of the charges (NaBr, LiBr) < (MgS, SrS) Then examine the ion sizes of each group and order by radius larger < smaller NaBr < LiBr < SrS < MgS

31 Lattice Energy and Heat of Hydration Hsolution,NaCl = Hhydration,NaCl(aq) UNaCl Hhydration,NaCl(aq) = Hhydration,Na+(g) + Hhydration,Cl (g)

32 Enthalpies of Hydration of Selected Cations and Anions

33 Hfinal ΔHsoln = 3.9kJ/mol NaCl

34 NaOH

35 NH 4 NO 3

36 Solubility of Gases in Water

37 Solubility of Gases in Water Solubility generally given in moles/l (M) Generally lower solubility than ionic or polar covalent solids Solubility decreases as temperature increases

38 Temperature Dependance of Gas Solubility

39 Henry s Law: The concentration of a sparingly soluble, chemically unreactive gas in a liquid is proportional to the partial pressure of the gas. C gas =k H P gas C = concentration of the gas in solution k H = Henry s law constant P gas = partial pressure of gas

40 Henry s Law: C gas =k H P gas

41 Henry s Law Constants for Gas Solubility in Water at 20ºC:

42 Pressure Dependance of Gas Solubility

43 Pressure Dependance of Gas Solubility