Chapter 9

Lewis Theory of Chemical Bonding

Lewis Bonding Theory

Emphasizes valence electrons to explain bonding

Lewis structures - Electron Dot Structures

Lewis structures allow us to predict many properties of molecules - molecular stability, shape, size, polarity

Why Do Atoms Bond?

A chemical bond forms when the potential energy of the bonded atoms is less than the potential energy of

the separate atoms.

To calculate this potential energy, you need to consider the following interactions:

nucleus–to–nucleus repulsions

electron–to–electron repulsions

nucleus–to–electron attractions

Types of Bonds

Types of Atoms Type of Bond Bond Characteristic

metals to nonmetals Ionic electrons

transferrednonmetals to

nonmetals Covalent electrons shared

metals to metals Metallic electrons pooled

Types of Bonding

Ionic Bonds

Metal atoms lose an electrons and become

cations.

Nonmetal atoms gain electrons and become

anions.

The oppositely charged ions are then attracted to each other, resulting in an

ionic bond.

Covalent BondsNonmetal atoms have relatively high

ionization energies, so it is difficult to remove electrons from them.

When nonmetals bond together, it is better in terms of potential energy for the atoms to share valence electrons.

The potential energy is lowest when the

electrons are between the nuclei.

Atoms held together because shared electrons are attracted to both nuclei.

Metallic BondsThe relatively low

ionization energy of metals allows them to lose electrons

easily.

Metal atoms release their valence electrons to be

shared as a pool by all the atoms/ions in the metal.

An organization of metal cation islands in a sea of

electrons.

Bonding results from attraction of cation for

the delocalized electrons.

Valence Electrons & Bonding

Because valence electrons are held most loosely, and

Because chemical bonding involves the transfer or sharing of electrons between two or more atoms,

Valence electrons are the most important in

bonding.

Determining the Number of Valence Electrons in an Atom

The column number on the Periodic Table will tell you how many valence electrons a main group atom has.

IA IIA IIIA IVA VA VIA VIIA VIIIA

Li Be B C N O F Ne

1e- 2e- 3e- 4e- 5e- 6e- 7e- 8e-

Lewis Structures of AtomsWe represent the valence electrons of main-group elements

as dots surrounding the symbol for the element.

H

Li Be B C N O F Ne

He

Na Mg Al Si P S Cl Ar

B C N

P

IA

IIA IIIA

IVA VA VIA VIIA

VIIIA

Practice – Write the Lewis structure for

Arsenic

A lithium ion A fluoride ion

•••

••

As

Stable Electron Arrangementsand Ion Charge

Metals form cations by losing enough electrons to get the same electron configuration as the previous noble gas.

Nonmetals form anions by gaining enough electrons to get the same electron configuration as the next noble gas.

The noble gas electron configuration must be very stable.

Main-group ions and the noble gas configurations.

Lewis Bonding Theory ⇒ Octet Rule

When atoms bond, they tend to gain, lose, or share electrons to result in eight valence electrons

ns2np6 (noble gas configuration)

ExceptionsH, Li, Be, B attain an electron configuration like He

He = two valence electrons (a duet)Li loses its one valence electron H may share or gain one electron

It commonly loses its one electron to become H+ Be loses two electrons to become Be2+

It commonly shares its two electrons in covalent bonds, resulting in four valence electrons

B loses three electrons to become B3+

It commonly shares its three electrons in covalent bonds, resulting in six valence electrons

Expanded octets for elements in Period 3 or below

Lewis Theory and Ionic Bonding

Lewis symbols can be used to represent the transfer of electrons from metal atom to nonmetal atom, resulting in ions that are attracted to each other and therefore bond.

Sodium Chloride Formation

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.

Predicting Ionic FormulasUsing Lewis Symbols

Electrons are transferred until the metal loses all its valence electrons and the nonmetal has an octet.

Numbers of atoms are adjusted so the electron transfer comes out even.

Li2O

Ca · · Cl · · ·

· · · ·

Use Lewis theory to predict the chemical formula of calcium chloride

Cl · · ·

· · · · CaCl2

Ca2+

Use Lewis symbols to predict the formula of an ionic compound made from reacting a metal, M, that has two valence

electrons with a nonmetal, X, that has five valence electrons

M3X23M2+

2X3-

Sr3N2

Energetics of Ionic Bond FormationThe ionization energy of the metal is endothermic:

Na(s) → Na+(g) + 1 e ─ " ΔH° = +496 kJ/mol

The electron affinity of the 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?

Na(s) + ½Cl2(g) → NaCl(s)" ΔH°f = + “something”

Ionic Bonding & the Crystal LatticeThe 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.

Crystal Lattice

Electrostatic attraction is nondirectional!!

There is no direct anion–cation pair “bond”

Therefore, there is no ionic molecule.

The chemical formula for an ionic compound is an empirical formula, simply giving the ratio of ions based on charge balance.

Lattice EnergyThe 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.

Practice – Given the information below, determine the lattice energy of NaCl

Na(s) → Na(g) +108 kJ½ Cl2(g) → Cl(g) +½(244 kJ)Na(g) → Na+(g) +496 kJCl (g) → Cl−(g) −349 kJNa(s) + ½ Cl2(g) → NaCl(s) −411 kJ

Na+ (g) + Cl−(g) → NaCl(s) ΔH (NaCl lattice)?

Determining Lattice EnergyThe Born–Haber Cycle

The Born–Haber Cycle is a hypothetical series of reactions that represents the formation of an ionic compound from its constituent elements.

The reactions are chosen so that the change in

enthalpy of each reaction is known except for the last one, which is the lattice energy.

Naº (s) + ½ Cl2 (g) NaCl (s)

Born–Haber Cyclefor NaCl

ΔH°f (metal atoms, g)

ΔH°f (nonmetal atoms, g)

ΔH°f (cations, g)

ΔH°f (anions, g)

ΔH°(crystal lattice)

ΔH°f (salt)

separating atoms

forming ions

forming lattice

Born–Haber Cycle

Use Hess’s Law to add up enthalpy changes of other reactions to determine the lattice energy.

ΔH°f(salt) = ΔH°f(metal atoms, g) + ΔH°f(nonmetal atoms, g)

+ ΔH°f(cations, g) + ΔH°f(anions, g)

+ ΔH°(crystal lattice)

ΔH°f(NaCl, s) = ΔH°f [Na(s)--->Na(g)] (Heat of vaporization) + ΔH°f (Cl–Cl bond energy) + Na 1st Ionization Energy + Cl Electron Affinity

+ NaCl Lattice Energy

ΔH°f(NaCl, s) =

ΔH°f(Na atoms,g) + ΔH°f(Cl atoms,g)

+ ΔH°f(Na+,g) + ΔH°f(Cl−,g)

+ ΔH°(NaCl lattice)

Na(s) → Na(g) +108 kJ½ Cl2(g) → Cl(g) +½(244 kJ)Na(g) → Na+(g) +496 kJCl (g) → Cl−(g) −349 kJNa+ (g) + Cl−(g) → NaCl(s) ΔH (NaCl lattice)Na(s) + ½ Cl2(g) → NaCl(s) −411 kJ (measured in an experiment)

NaCl Lattice Energy = (−411 kJ) − [(+108 kJ) + (+122 kJ) + (+496 kJ) + (−349 kJ) ]

= −788 kJ

NaCl Lattice Energy = ΔH°f(NaCl, s) − [ΔH°f(Na atoms,g) + ΔH°f(Cl–Cl bond energy) + Na 1st Ionization Energy + Cl Electron Affinity]

Practice – Given the information below, determine the lattice energy of MgCl2

Mg(s) ➔ Mg(g) ΔH1°f = +147.1 kJ/mol½ Cl2(g) ➔ Cl(g) ΔH2°f = +122 kJ/molMg(g) ➔ Mg+(g) ΔH3°f = +738 kJ/molMg+(g) ➔ Mg2+(g) ΔH4°f = +1450 kJ/molCl(g) ➔ Cl−(g) ΔH5°f = −349 kJ/molMg(s) + Cl2(g) ➔ MgCl2(s) ΔH6°f = −641 kJ/mol

Practice – Given the information below, determine the lattice energy of MgCl2

Mg(s) ➔ Mg(g) ΔH1°f = +147.1 kJ/mol2{½ Cl2(g) ➔ Cl(g)} 2ΔH2°f = 2(+122 kJ/mol)Mg(g) ➔ Mg+(g) ΔH3°f = +738 kJ/molMg+(g) ➔ Mg2+(g) ΔH4°f = +1450 kJ/mol2{Cl(g) ➔ Cl−(g)} 2ΔH5°f = 2(−349 kJ/mol)Mg2+(g) + 2 Cl−(g) ➔ MgCl2(s) ΔH° lattice energy = ? kJ/mol

Mg(s) + Cl2(g) ➔ MgCl2(s) ΔH6°f = −641 kJ/mol

Trends in Lattice EnergyIon Size

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 attractionweaker attraction = smaller lattice energy

Lattice Energy vs. Ion Size

Trends in Lattice EnergyIon 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 attractionstronger attraction = larger lattice energy

Of the two factors, ion charge is generally more important

Lattice Energies of Some Ionic Solids (kJ/mole)(M+ + X- → MX)

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 4,443

Mg2+ 2,957 2,524 2,440 2,327 3,791

Ca2+ 2,630 2,258 2,176 2,074 3,401

Al3+ 5,215 5,492 5,361 5,218 15,916

Anions

(2M+ + X2- → M2X)(M2+ + 2X- → MX2)(M2+ + X2- → MX)(M3+ + X- → MX3)(2M3+ + 3X- → M2X3)

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

Ionic Bonding-Model vs. ObservationsLewis theory implies strong attractions between ions.

Lewis theory predicts high melting points and boiling points for ionic compounds.

The stronger the attraction (larger the lattice energy),

the higher the melting point.

Ionic compounds have high melting points and boiling points (MP generally > 300 °C).

All ionic compounds are solids at room temperature.

Melting an ionic solid

Ionic Compounds Melt

Lewis theory implies that the positions of the ions in the crystal lattice are critical to the stability of the structure

Lewis theory predicts that moving ions out of

position should therefore be difficult, and ionic solids should be hard

Ionic solids are relatively hard(compared to most molecular solids)

Ionic Bonding-Model vs. Observations

Lewis theory implies that if the ions are displaced from their position in the crystal lattice, that repulsive forces should occur

This predicts the crystal will become unstable and break apart. Lewis theory predicts ionic solids will be brittle.

Ionic solids are brittle. When struck they shatter.

Ionic Bonding-Model vs. Observations

To conduct electricity, a material must have charged particles that are able to flow through the material

Lewis theory implies that, in the ionic solid, the ions are locked in position and cannot move around

Lewis theory predicts that ionic solids should not

conduct electricity

Ionic solids do not conduct electricity

Ionic Bonding-Model vs. Observations

Conductivity of NaCl

in NaCl(s), the

ions are stuck in

position and not

allowed to move

to the charged

rods

Lewis theory implies that, in the liquid state or when dissolved in water, the ions will have the ability to move around

Lewis theory predicts that both a liquid ionic compound

and an ionic compound dissolved in water should conduct electricity

Ionic compounds conduct electricity in the liquid state or when dissolved in water

Ionic Bonding-Model vs. Observations

Conductivity of NaCl

in NaCl(aq),

the ions are

separated and

allowed to

move to the

charged rods