1 chapter 9 models of chemical bonding -...

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© 2006 Brooks/Cole - Thomson

Chapter 9 MODELS OF CHEMICAL BONDING

HB+ H

AH

BH

A

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A comparison of metals and nonmetals

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Types of Chemical Bonding

Ionic bonding involves the transfer of

electrons and is usually observed when a

metal bonds to a nonmetal.

Covalent bonding involves the

sharing of electrons and is usually

observed when a nonmetal bonds to a

nonmetal.

Metallic bonding involves electron

pooling and occurs when a metal bonds

to another metal.

9.1 Atomic Properties & Chemical Bonds

Chemical bond: A force that holds atoms together

in a molecule or compound

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Three models of chemical bonding

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Gradations in bond type among Period 3

and Group 4A elements

Most bonds are somewhere in between ionic and covalent.

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Lewis Symbols & the Octet Rules

Lewis symbols for atoms show

valence electrons only

Rules of the Game

# of valence electrons of a main group atom = Group number

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Lewis Symbols & The Octet Rule

• The OCTET RULE: Many atoms gain or lose electrons so as to end up with the same # of electrons as the noble gas closest to them on the periodic table

Example: Phosphorus has 5 valence electrons and want to get 3 more electrons in its valence shell

P Group 5A(15) 1s2 2s2 2p6 3s2 3p3

Br Group 7A(17) [Ar] 4s2 3d10 4p5

Core = [Ar] 3d10 , valence = 4s2 4p5

Bromine has 7 valence electrons and want to

get 1 more electron in its valence shell

Why do compounds form? -Elements with unfilled valence shells usually form compounds with other elements to gain stability.

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–Usually made up of a metal ion and one or more nonmetal ions

»Metals have low IE and tend to lose electrons (to form cations) – e.g., Li becomes Li+ ion.

»Nonmetals have large negative values of EA and tend to gain electrons (to form anions) – e.g., F becomes F- ion.

9.2 The Ionic Bonding Model

- Why don’t

compounds such

as Li2F, LiF2

exist?

Li • F ••

• •• •• Li+ + F

••

•• ••

••

-

Li 1s22s1 + F 1s22p5 → Li+ 1s2 + F- 1s22s22p6

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The importance of Lattice Energy DH0lattice

Why do ionic compounds form?

The Born-Haber Cycle

to determine Lattice

Energy

Lattice energy is the energy required to separate 1 mol of

an ionic solid into gaseous ions.

Lattice energy is a measure of the strength of the ionic

bond.

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Periodic Trends in Lattice Energy

Coulomb’s Law:

- As ionic size increases,

lattice energy decreases.

Lattice energy therefore

decreases down a group on

the periodic table.

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© 2006 Brooks/Cole - Thomson 11

Periodic Trends in Lattice Energy

NaCl, Na+ and Cl-,

m.p. 804 oC

MgO, Mg2+ and O2-

m.p. 2800 oC

Coulomb’s Law:

- As ionic

charge increases,

lattice energy

increases.

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Properties of Ionic Compounds

• Ionic compounds tend to be hard, rigid, and brittle, with high melting points.

• Ionic compounds do not conduct electricity in the solid state.

– In the solid state, the ions are fixed in place in the lattice and do not move.

• Ionic compounds conduct electricity when melted or dissolved.

– In the liquid state or in solution, the ions are free to move and carry a current.

ionic

compounds are

easily cracked

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Interionic attractions are so

strong that when an ionic

compound is vaporized, ion

pairs are formed.

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9.3 The Covalent Bonding Model

Covalent bond arises from the mutual attraction of 2

nuclei for the same electrons. Electron sharing

results.

HB+ H

AH

BH

A

Covalent bond is a balance of attractive and repulsive

forces.

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The Covalent Bonding Model

Formation of a

covalent

bond results

in greater

electron

density

between the

nuclei.

Covalent bond

formation in H2.

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Bonding Pairs & Lone Pairs of electrons

• Valence electrons are distributed as

shared or BONDING PAIRS and

unshared or LONE PAIRS.

This is called a LEWIS ELECTRON DOT

structure of a molecule.

G. N. Lewis

1875 - 1946

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Covalent Bond Properties: Order, Energy. & Length

• What is the effect of bonding and structure on molecular properties?

Free rotation

around C–C single

bond

No rotation around

C=C double bond

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Bond Order # of electron pairs being shared between

a given pair of atoms

Double bond Single bond

Triple

bond

Acrylonitrile

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• —measured by the energy required to break a covalent bond. Bond Enthalpy has positive values.

• BOND STRENGTH (kJ/mol)

H—H 436

C—C 346

C=C 610

CC 835

NN 945

The GREATER the bond order the HIGHER the

bond strength and the SHORTER the bond length.

Bond Energy (Bond Strength, Bond Enthalpy)

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Bond Order Length Strength

HO—OH

O=O

1 142 pm 210 kJ/mol

2 121 498 kJ/mol

1.5 128 ?

Bond Strength

O O•••••

••

••

••O

••

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Bond Length

Bond length depends on size of bonded atoms.

H—F

H—Cl

H—I

Bond distances

measured in

Angstrom units

where 1 A = 10-2 pm =

10-8 cm

The distance between the

nuclei of two bonded atoms.

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Bond length also depends on bond order.

Bond distances measured in

Angstrom units where 1 A = 10-2 pm.

Bond Length

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Physical Properties of Covalent Substances

1) Poor electrical conductivity

2) Physical properties of

- Molecular covalent substances: gas, liquid, low-

melting point solid due to Strong and localized intramolecular (bonding) forces

Weak intermolecular forces

- Network covalent solids:

Covalent bonds

throughout the sample

Quartz (SiO2) mp 1550 0C

Diamond mp 3550 0C

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9.4 Bond Energy & Chemical Change

∆H0rxn = (Sum of energy required to break bonds)

+ (Sum of energy generated by newly formed bonds)

Using Bond Energies to Calculate DH0rxn

Guide: 1) Break ALL the reactant bonds to obtain individual atoms

2) Use the atoms to form ALL the product bonds

3) Add the bond energies with appropriate signs to obtain DH0rxn

∆H0rxn > 0 for broken bonds

∆H0rxn < 0 for formed bonds

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Problem: Using bond energies to calculate DH°rxn for HF

formation. H2 (g) + F2 (g) 2 HF (g)

∆H0rxn = (BEH2

+ BEF2) + 2 BEHF

∆H0rxn = (432 kJ/mol + 159 kJ/mol) + (– 2x565 kJ/mol)

∆H0rxn = -539 kJ/mol

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Using bond energies to calculate DH°rxn for the combustion of

methane. CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)

∆H0rxn = [4 BECH4

+ 2 BEO=O] + [2 BEC=O + 4 BEO-H]

∆H0rxn = [(4 x 413) + (2 x 498)] + [(2 x -799) + (4 x -467)]

∆H0rxn = -818 kJ/mol

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Relative bond

strength and

energy from

fuels.

Bond Strengths &

Heat Released from Fuels & Foods

Weaker bonds

such as C-H

bonds (less

stable, more

reactive) are

easier to break

than stronger

bonds such as

C-O bonds

(more stable,

less reactive)

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Bond Strengths &

Heat Released from Fuels & Foods

Fuels with

fewer bonds

to O release

more energy.

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9.5 Electronegativity and Bond Polarity

A covalent bond in which the shared

electron pair is not shared equally, but

remains closer to one atom than the other,

is a polar covalent bond.

Unequal sharing of electrons causes the

more electronegative atom of the bond

to be partially negative and the less

electronegative atom to be partially

positive.

The ability of an atom in a covalent bond to

attract the shared electron pair is called its

electronegativity.

H Cl••

••

+ -

••

Due to the bond polarity, the H—Cl bond energy is GREATER than expected for a “pure” covalent bond.

BOND ENERGY

“pure” bond BEHCl = 339 kJ/mol

real bond BEHCl = 432 kJ/mol

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A measure of the ability of a bonded atom to attract shared electrons

Relative values of EN determine BOND POLARITY

Electronegativity (EN)

F has

highest

EN

value

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© 2006 Brooks/Cole - Thomson Figure 9.22

Electronegativity and atomic size.

Trends in Electronegativity

EN increases

as atomic size

decreases.

Nonmetals

have higher

EN than

metals.

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Electronegativity and Oxidation Number

• The more electronegative atom is assigned all the shared electrons.

• The less electronegative atom is assigned none of the shared electrons.

• Each atom in a bond is assigned all of its unshared electrons.

Electronegativities can be used to assign oxidation numbers (ON):

O.N. = # of valence e- - (# of shared e- + # of unshared e-)

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Example:

Cl is more electronegative than H, so for Cl:

valence e- = 7

shared e- = 2

unshared e- = 6

O.N. = 7 – (2 + 6) = -1

H is less electronegative than Cl, so for H:

valence e- = 1

shared e- = 0 (all shared e- assigned to Cl)

unshared e- = 0

O.N. = 1 – (0 + 0) = +1

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Depicting Polar Bonds

The unequal sharing of electrons can be depicted by a polar

arrow. The head of the arrow points to the more electronegative

element.

A polar bond can also be marked using δ+ and δ- symbols.

Bond Polarity & Partial Ionic Character

H Cl••

••

+ -

••

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The Importance of DEN

- Percent ionic

character of a

bond increases

with DEN .

- No bonds are

purely ionic or

covalent .

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Determine Bond Polarity from DEN

Which bond is more polar?

O—H O—F

DEN 3.5 - 2.1 3.5 - 4.0

DEN 1.4 0.5

OH bond is more polar than OF bond, and direction of polarity is opposite in these cases.

O H

+-

O F+ -

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Properties

of the

Period 3

chlorides.

As DEN decreases, melting point and electrical conductivity

decrease because the bond type changes from ionic to polar

covalent to nonpolar covalent.

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

The Graduation in Bonding across a Period

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9.6 Introduction to Metallic Bonding

The Electron Sea Model: • All metal atoms in the sample contribute their valence

electrons to form a delocalized electron “sea”. • The metal “ions” (nuclei with core electrons) lie in an

orderly array within this mobile sea.

• All the atoms in the sample share the electrons.

• The metal is held together by the attraction between the metal “cations” and the “sea” of valence electrons.

Group

2A/2

Group

1A/1

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Properties of Metals • Metals are generally solids with moderate to high melting

points and much higher boiling points.

–Melting points decrease down a group and increase across a period.

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Properties of Metals - Mechanical properties: Metals can be shaped without breaking.

– The electron sea allows the metal ions to slide past each other.

Metals dent or

bend rather than

crack

- Electric Conductivity: Metals are good conductors of

electricity in both the solid and liquid states. The electron sea is mobile in both phases.

- Heat Conductivity: Metals are good conductors of heat.

1 chapter 9 models of chemical bonding -...

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