BONDING

TOPIC 4

Terms

Covalent Bonding• Bonds– Breaking them takes energy– Making them gives off energy

• Exothermic– More energy is given off than put in

• Endothermic – More energy is absorbed than given off

• Intramolecular Forces– Forces within molecules (ionic, covalent and metallic)

• Intermolecular Forces (IMF)– Forces between particles

Metal: K Non-Metal: Cl

Ionic Bonding

• If the electronegative difference between the atoms involved is =>1.8– There are always exceptions to this rule!

• Will conduct electricity in its molten or aqueous state (This test proves ionic)

Less e- = Less e- repulsion More e- = e- more repulsion.

+ -

Intramolecular Forces

Ionic Bonding• Just use the valence shell

• Be sure to include square brackets and charge after electron exchange.

Drawing Ionic Bonding

Na Cl

XElectrons are

in pairs

+

Special Note: The ionic bond is the electrostatic attraction between oppositely charged ions!

-Lewis Dot Diagram

Lewis Diagrams

Lewis Dot diagrams us the atoms valance shell electrons

Combine

C

Cl

Be

Br

Al

F

Fe

Cl

Mg

O

Intramolecular Forces

NaCl

• When in molten or aqueous state, ionic substances WILL conduct electricity, by the movement of (+) and (-) ions.

• This is different from how METALS conduct electricity!

Decomposition

2Na+(aq) + 2Cl-

(aq) 2Na(s) + Cl2(g)CA

THO

DE

(-)

ANO

DE

(+)

-

++

+

++

+

-

--

-+ +

+ +

+ +

Conductivity is FINITE

Metal: K Non-Metal: Cl

Ionic Compounds

+ - + - +-

-+

+ - ++

--

-+-+- +

+ +-+-

• No bonds are made!!!

• Static attractions holds them together. (opposites attract)

• When a force is applied, ionic compounds will make a clean break.

Force

Like charges repel

Giant Ionic Lattices

• Physical characteristics• Hard and brittle• Solid doesn’t conduct Electricity• More soluble in water than other solvents• High MP and BP

CationAnion

Table SaltNaCl

Cubic or IsometricGiant Ionic Lattices

CassiteriteSnO2

Tetragonal

Giant Ionic Lattices

AgagoniteCaCO3

Orthorhombic

• Also found in mollusk shells and coral

Giant Ionic Lattices

BerylBe3Al2(SiO 3)6

Hexagonal

Giant Ionic Lattices

QuartzSiO2

Trigonal

Giant Ionic Lattices

BerylBe3Al2(SiO 3)6

Ionic BondingGiant Ionic Lattices

Copper(II) SulfateCuSO4

Triclinic

Giant Ionic Lattices

Intramolecular Forces

Multiple Ions• Transition metals can have

multiple ions.• Ones you should know.

Transition Metals

Cu2+Cu+Fe2+ Fe3+

Copper(I) Oxide Iron (III) OxideIron(II) Oxide Copper(II) Oxide

Ions

Polyatomic Ions• Be sure to review your polyatomic

ions!!!

Reminder

NO3- SO4

-2 PO4-3

OH- CO3-2

NH4+

HCO3-

Covalent Bonding

Topic 4

Intramolecular Forces

COVALENT BONDING

• If the electronegative difference between the atoms involved is <1.8

• Will NOT conduct electricity• Electrons are shared

Covalent Bonding

X

H Cl

3.02.1

Differences |3-2.1|

=0.9

Special Note: The covalent bond is the electrostatic attraction between pairs of e- and positively charged nuclei!

Questions

• For ionic compounds to form the valance shells of both metal and non-metal must be full!!

Na Ca

Li Na

+

+

+

+

Cl

O

CO3

SO3

• What is the chemical formula? • What is the names for each?

K + NO3

Review

Intramolecular Forces

COVALENT BONDING

• Structural formula

• Lewis structure

Covalent Bonding

H C

H

H

H Cl

X

XX

X

H

H

H

H

Intramolecular Forces

COVALENT

BONDING

• 1) Sum all valence e-

• 2) Subtract 2e- for every bond• 3) Place e- around periphery atoms to

form octets. The remaining around central atom

• 4) All atoms MUST be paired!!!!!!

Lewis StructuresH2O

1

1

6

8

H

H

O

- 4 = 4

Hydrogen can only hold 2e- remaining must be paired on

Oxygen

Intramolecular Forces

Covalent Bonding

• Draw the following Lewis structures

• H2 Cl2• O2 N2

• HCN C2H6

• C2H4 C2H2

Lewis Structures

HL: PCl5, PCl4+, PCl6

-

and XeF4

Intramolecular Forces

Covalent Bonding

• Coordinate or dative covalent bonds• When both e- are shared from the same

atom. (Not one from each as before)• Occurs when a non bonding e- pair

donates an e- to an e- deficient atom.

Special Lewis Structures

H

H

NH

+

Electrophile

H

Lone pair of e-

+

Intramolecular Forces

Covalent Bonding

• Draw the following Lewis structures• CO

• H3O+

Special Lewis Structures

Intramolecular Forces

CO32-

• More bonds = more strength & shorter bonds

• Resonance structures – Bond length is longer than a double bond but

shorter than a single bond

Length, Strength & Hybrid Resonance

CO O

O2-

CO O

O2-

CO O

O2-

Don’t forget to show the e- pairs!!

Intramolecular Forces

CO32-

• Compare the two molecules

• Ethyne has stronger and shorter bonds• • C=O bond is stronger and shorter due to

Oxygen being more electronegative

Length & Strength

Ethene

R = Functional Group

CR OH

OC

H

H

C

H

H

Carboxylic Acid

C HCH

Ethyne

Intramolecular Forces

Covalent

Bonding

• Non-Metals are fighting for e-

• Atom with larger electronegativity will hold the e- closer to itself.

• Atoms become slightly charged.

Bond Polarity

ClXH

δ-δ+

Dipole Moment

Intramolecular Forces

Covalent

Bonding

• BF3

• Actual structure: Boron is e- deficient• This is known because of its reactivity towards

electron rich molecules such as NH3

• CNOF all obey the octet rule.

Exceptions to the Octet Rule

FF

B

F

• SO42-

• Single bonds (8 e- around S)• Double bonds (12 e- around S)

• Formal Charge = (# valence e- on free atom) – (# valence e- assigned to the atom in the molecule)

• (Valence e-)assigned = (# lone pair e-) + ½ (# of shared e-)

• 1) Molecules attempt to achieve Formal Charge as close to 0 as possible.

• 2) Any negative Formal charge will reside on most electronegative atom.

Formal Charge

Covalent Bonding

Intramolecular Forces

Intramolecular Forces

Covalent

Bonding

• VSEPR (Valence Shell Electron Pair Repulsion)

• Paired e- attempt to get as far away from each other as possible.

• Multiple bonds still count only as 1 pair!!

VSEPR (shape)

OO C

FF

B

F

3 Pairs of e-

120o

2 Pairs of e-

180o

OO

C

O2-

Intramolecular Forces

Covalent

Bonding

• Tetrahedral

• Lone pair e- have increased charge density and require more room

• More repulsion from lone pair will decrease bond angle.

VSEPR

H

H

C

H

H

4 Pairs of e-

109.5o

H

H

N

H

Lone pair107o

H

O

H

Lone pair104.5o

Intramolecular Forces

Covalent Bonding

• Predict the shape AND bond angles

• H2S PbCl4 H2CO SO2

• NO3- PH3 NO2

-

• NH2- POCl3 CO2

Home Work

HL VSEPRMolecule Shape Total valance

electrons Bond Pairs Non Bonding

Electron pairAngle

BeF2 Linear 180

BeF3 Triangular Planar

120

SO2 Bent 117

CH4 Tetrahedral 109.5

NH3 Trigonal pyramidal

107

H2O Bent 104

HL VSEPRMolecule Shape Total

Valance electrons

Bond Pairs Non Bonding Electron pair

Angle

PCl5 Triangular Bipyramidal

90 & 120

SF4 Seesaw 90 & ≈117

T-Shape 90

CF6 Octahedral 180

IF5 Square Pyramidal

90

XeF4 Square Planar

≈88

Expanded Valance Shell (14.1)

• Molecules with more than 8 electrons• Electron promotion:

Dipole Moment

Covalent Bonding

• Polarity effects state change (physical change)

• Unequal sharing causes a dipole moment to form

• Q: Why is BF3 non-polar whereas PF3 is polar?

Molecule Polarity (4.2.6)

δ+

2δ-

δ+H

O

H

ClH δ-

δ+

Non Polar

H

H

C

H

δ-

δ+ H

H

C

Cl

Cl

Hybridization (14.2.2)• Sigma bond: σ (single bond)

– Axial overlap of orbital’s

Cl

1s1

H

2px2 py

2 pz2

Hybridization (14.2)• Sigma bond: σ (single bond)

– Axial overlap of orbital’s

ClCl

Hybridization (14.2)• Pi bond: π(Double bond, one σ bond)

– Parallel overlap of orbital’s

NN OO

Hybridization (14.2.3)• Hybridization electron promotion

– New Orbital sp3

2px2 py

2 pz2

2s2Ground

StateC

Excited State

4 Equal orbital`s capable of holding a maximum of 2 electrons each

Hybridization (14.2)

• How to determine Hybridized orbital`s– Look at the shape

Shape High Electron dense regions

Hybridized Orbital

sp 2 Tetrahedral

Sp2 3 Trigonal planar

sp3 4 Linear

sp3d 5 Trigonal bi-pyramidal

sp3d2 6 Octahedral (Square bi-pyramidal)

Allotropes

Giant Covalent

• 1) Diamond (Tetrahedron, localized e-)– Very hard and does not conduct electricity

• 2) Fullerenes (C60) Hexagonal and pentagonal rings– Nanotubes

Carbon

C

CC

C

C

Allotropes

Giant Covalent

• 3) Graphite (Planar, delocalized e-)– Weak pi bonding between sheets cause it to

conduct electricity and be slippery.– Bonds are shorter than a tetrahedral due to the

pi bonding

Carbon

C

C

CC

C

C

HL: sp hybridDelocalized electrons

able to move

Weak Pi Bonds

C6H6• Planar, delocalized e-

– Regular bonding would predict an alternating double bond (Resonance structure)

– Hybrid theory shows sp2 configuration

Benzene (14.3)

C

C

CC

C

C

Pi bonds overlap allowing for electrons to be delocalized over the entire molecule.

Intramolecular Forces

SiliconTetrahedron Configuration

Similar to diamond

Silicon

Si

Si

Si

Si

Si

Si

Si

SiSi

Si

Intramolecular Forces

Quartz

• Single bonds formed between Oxygen to satisfy the octet.

• HL: Less overlap in the P-sub orbital due to atomic size difference therefore Pi bonds do not form.

Silicon & Silicon dioxide

SiOO

O

O

SiO2 but based on a network

of SiO4

Metallic Bonding

Topic 4

Intramolecular Forces

Metallic Bonding

• In solid state

• Outer e- are delocalized and free to move about

• Bond is a result of electrostatic attraction between Fixed positive metal ions and delocalized e-

Metallic Bonding+ +

+

+

+

+ +

+

+

++

++

+

+ ++-

-

-

- --

Conductivity is INFINITE

-

-

Sea of electrons

Metallic Bonding

• The ability for a material to be pounded into thin sheets.

• Aluminum Foil

• Swords and Folding

+ +

+

+

+

+ +

+

+

++

++

+

+ ++

-

-

-

- --

-

-

MalleabilityPhysical Properties

Ductility

Metallic Bonding

• The ability for a material to be pulled into wire

• Or in this case extruded into a wire

Physical Properties

+

+

+ +

+

+

+

+

++

+

+

Electrons have been excluded

Metallic Bonding

• Because e- can move easily it can conduct energy. (Heat or electricity)

• MP related to attractive force (between atoms) • 1) Size of Cation(+)

• 2) # of valence e-

• 3) Atom packing

• Size increases MP decreases: • Giant Covalent substances have very high

mp

Physical Properties

Metallic Bonding

• Same element but different structure• Carbon

• Diamond• Graphite• Fulluron

Allotrops

INTERMOLECULAR FORCES

Topic 4

Intermolecular Forces

IMF• Van der Waals Forces

van der Waals’ ForcesIntermolecular Forces (4.3.1)

Charge Induction Charge Induction

d+

d- d+

d- d-d+

Intermolecular Forces

IMF• Polar molecules (polar covalent) have

slightly charged ends• Opposites attract.• Large electronegative difference =

stronger attraction.

Dipole-Dipole (4.3.1)

H

H

C

Cl

Cl

d-

d+

H

H

C

Cl

Cl

d-

d+

Intermolecular Forces

IMF• Hydrogen Bonding (F, O or N bonded to H)

• Due to small size and high electronegativity of non metals

• Creates a large charge difference • Basically a super strong dipole-dipole bond

van der Waals’ ForcesHydrogen Bonding (4.3.1)

d+d+ H

O

H

d-

d+

d+

H

OH

d-

Get a picture of group 4,5,6,7 boiling points for hydrides

Key question is why does water have an abnormally high BP?

H bonding with O, F and N

IMF• Phase change when IMF are overcome• Be sure to explain using the words IMF and

how they affect the bonds BETWEEN particles.• Van der Waals’ Forces are ALWAYS present!!!

Boiling Point Trends (4.3.2)Intermolecular

Forces

• Van der Waal’s: Lowest MP, Non polar• Butane (C4H10)

• Dipole-dipole: Slightly miscible• Propanone C3H6O

• Hydrogen Bonding: Miscible with polar substances• H2O

• Ionic Bonding: Only conducts electricity when liquid or aqueous. (Decomposition when it does)• NaCl

• Metallic Bonding: Conducts electricity, not water soluble, MP regulated by, valance, size and packing.• Fe

• Giant Covalent: Highest MP, Insoluble in both non-polar and polar solvents. Does not conduct electricity except for graphite.• Diamond and Graphite (Allotropes)

Physical Properties

Increasing Melting Point

Bonding Questions

• Compare the following for B.P

• HF and HCl• H2O and H2S

• NH3 and PH3

• CH3OCH3 and CH3CH2OH

• CH3CH2CH3, CH3CHO and CH3CH2OH

HL Material

Hybridization (14.2)

• Sigma bond: σ (single bond)– Axial overlap of orbital’s

Hybridization (14.2)

• Sigma bond: σ (single bond)– Axial overlap of orbital’s

Hybridization (14.2)

• Sigma bond: σ (single bond)– Axial overlap of orbital’s

Hybridization (14.2)

• Sigma bond: σ (single bond)– Axial overlap of orbital’s

Lattice Formation

• Where the heat comes from

• Route 1: A + B + C + E• Route 2: F

• Hess’s law: A + B +C + E = F+107 + 122 + 496 + (-349) + E = -411E = -787 kJ mol-1

Intramolecular Forces

NaCl

• 1) Production of Gaseous atoms• 2) Formation of Gaseous ions• 3) Production of solid ionic lattice

Lattice Enthalpy

Na(s) + ½ Cl2(g) Na+Cl- or NaCl

1) Na(s) Na(g) ½ Cl2(g) Cl(g)

3) Na+(g) + Cl-

(g) NaCl(s)

2) Na(g) Na+(g) + e- Cl(g) + e- Cl- (g)

Born-Haber CycleNa+

(g) + Cl-(g)

Na(s) + Cl2(g)

Na(g) Cl(g)

NaCl(s)

ΔHθI.E.

1st Ionization of Na+496 kJ mol-1

ΔHθat

Atomization of Na-107 kJ mol-1

ΔHθat

Atomization of Cl+122kJ mol-1

ΔHθE.A.

1st electron affinity of Cl-349 kJ mol-1

ΔHθf

Formation of NaCl-411 kJ mol-1

Lattice Enthalpy

BA

DC

E

FExothermicEndothermic

OO C

Spare Parts

O C

ClXH δ-

δ+

Cl

HH

C

H

N

H

H

NH

+H

-

H

O

H

H

H

C

H

H

C

C

CC

C

C