Upload
others
View
0
Download
0
Embed Size (px)
Citation preview
CH1810 Lecture #1
Solutions of Ionic Compounds
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.
Gas in gas Air (O2 in N2)
Gas in liquid Carbonated water (CO2 in H2O)
Gas in solid H2 in palladium metal
Liquid in liquid Gasoline, tequila
Liquid in solid Dental amalgam (Hg in Ag)
Solid in liquid Salt water (NaCl in H2O)
Solid in solid Sterling silver (Cu in Ag)
Kinds of Solutions
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.
Attractive Forces and Solubility
All Molecules
Polar Molecules
Molecules containing O-H, N-H, or F-H
Bonds
Dispersion forces
Dipole forces
H-bonding
Hierarchy of Intermolecular Forces
Charged particles
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, NH2, Cl
Nonpolar molecules dissolve in nonpolar solvents
hydrophobic groups = C-H, C-C
Immiscible Liquids
Pentane, C5H12, is a nonpolar molecule.
Water is a H2O, 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.
Nonpolar Solvents
Polar Solvents
Ethanol
(ethyl alcohol)
Water
Dichloromethane
(methylene chloride)
Practice – Choose the substance in each pair that is more soluble in water
a) CH3OH CH3CHF2
b) CH3CH2CH2CH2CH3 CH3Cl
Salt vs. Sugar Dissolved in Water
ionic compounds
dissociate into ions when
they dissolve
molecular compounds
do not dissociate when
they dissolve
Ionic compounds dissociate into ions when they
dissolve.
Molecular compounds do not dissociate into ions
when they dissolve.
• When ionic compounds dissolve in water, the
anions and cations are separated from each
other. This is called dissociation.
Na2S(aq) → 2 Na+(aq) + S2-(aq)
• When compounds containing polyatomic ions
dissociate, the polyatomic group stays together
as one ion
Na2SO4(aq) → 2 Na+(aq) + SO42−(aq)
• When strong acids dissolve in water, the
molecule ionizes into H+ and anions
H2SO4(aq) → 2 H+(aq) + SO42−(aq)
Dissociation vs IonizationWhen ionic compounds dissolve in water, the anions
and cations are separated from each other. This is called dissociation.
When compounds containing polyatomic ions dissociate, the polyatomic group stays together as one ion.
When strong acids dissolve in water, the molecule ionizes into H+ and anions.
• When ionic compounds dissolve in water, the
anions and cations are separated from each
other. This is called dissociation.
Na2S(aq) → 2 Na+(aq) + S2-(aq)
• When compounds containing polyatomic ions
dissociate, the polyatomic group stays together
as one ion
Na2SO4(aq) → 2 Na+(aq) + SO42−(aq)
• When strong acids dissolve in water, the
molecule ionizes into H+ and anions
H2SO4(aq) → 2 H+(aq) + SO42−(aq)
• When ionic compounds dissolve in water, the
anions and cations are separated from each
other. This is called dissociation.
Na2S(aq) → 2 Na+(aq) + S2-(aq)
• When compounds containing polyatomic ions
dissociate, the polyatomic group stays together
as one ion
Na2SO4(aq) → 2 Na+(aq) + SO42−(aq)
• When strong acids dissolve in water, the
molecule ionizes into H+ and anions
H2SO4(aq) → 2 H+(aq) + SO42−(aq)
Energy Changes
During Dissolution of Ionic Solids
Dissolution of Ionic Compounds
Temperature and Solubility of Ionic Compounds
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.
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
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?
Therefore the formation of the ionic compound should be endothermic.
-
- -
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.
Lattice Energy
Lattice Energy
Lattice energy depends directly on size of charges and inversely on
distance between ions.
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
U = k(Q1)(Q2) d
Lattice Energy vs. Ion Size
Born-Haber Cycle
sublimation
bond breaking
ionization
electron affinity
lattice energy
Born-Haber Cycle
1) Na(s) + ½Cl2(g) → Na(g) + ½Cl2(g) + Hsub
2) Na(g) + ½Cl2(g) → Na(g) + Cl (g) + ½HBE
3) Na(g) + Cl (g) → Na+(g) + Cl (g) +HIE1
4) Na+(g) + Cl (g) → Na+(g) + Cl- (g) –HEA
5) Na+(g) + Cl–(g) → NaCl(s) - U
Born-Haber Cycle
U = Hf – ½ HBE – HEA – Hsub – HIE1
(sublimation)
(bond energy)
(ionization)
(electron affinity)
(lattice energy)
++
++
= ΔHf
Trends in Lattice Energy Ion Charge
Lattice Energy =
−910 kJ/mol
Lattice Energy =
−3414 kJ/mol
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
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
Cations F- Cl- Br- I- O2-
Li+ 1036 853 807 757 2,925
Na+ 923 787 747 704 2,695
K+ 821 715 682 649 2,360
Be2+ 3,505 3,020 2,914 2,800 4,443Mg2+ 2,957 2,524 2,440 2,327 3,791Ca2+ 2,630 2,258 2,176 2,074 3,401
Al3+ 5,215 5,492 5,361 5,218 15,916
Lattice Energies of Some Ionic Solids (kJ/mole)
Anions
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
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
∆Hsolution,NaCl = ∆Hhydration,NaCl(aq) – UNaCl
∆Hhydration,NaCl(aq) = ∆Hhydration,Na+(g) + ∆Hhydration,Cl–(g)
Lattice Energy and Heat of Hydration
Enthalpies of Hydration of Selected Cations and Anions
Dissolving ionic compounds in water.
NaCl
NaOH
NH4NO3
Hfinal
ΔHsoln = 3.9kJ/mol
Dissolving ionic compounds in water.
NaCl
NaOH
NH4NO3
Dissolving ionic compounds in water.
NaCl
NaOH
NH4NO3
Dissolving ionic compounds in water.
NaCl
NaOH
NH4NO3
Solubility of Gases in Water
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
Temperature Dependance of Gas Solubility
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 = concentration of the gas in solution kH = Henry’s law constant Pgas = partial pressure of gas
Cgas =kHPgas
Henry’s Law:
Cgas =kHPgas
Henry’s Law Constants for Gas Solubility in Water at 20ºC:
Pressure Dependance of Gas Solubility
Pressure Dependance of Gas Solubility