CH1810 Lecture #1 Solutions of Ionic...

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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

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)

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.

as one ion

H2SO4(aq) → 2 H+(aq) + SO42−(aq)

as one ion

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 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.

NH4NO3

Hfinal

ΔHsoln = 3.9kJ/mol

NH4NO3

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

CH1810 Lecture #1 Solutions of Ionic...

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