Chapter 7Chapter 7Chapter 7Chapter 7

11

Covalent BondsCovalent Bonds

and Molecularand Molecular

StructureStructure

Comparing Ionic and Covalent Bonds

Ionic bonds

must be broken

Intermolecular

forces (much weaker than

bonds) must be

broken

Ionic Bonds• Ionic bonds are very strong, so separating ions requires a lot of energy

– high melting points

– high boiling points

• Crystals are hard and brittle

– Crystal lattice is an arrangement of ions of opposite charge surrounding one another in three dimensions

• Electrical insulators when solid, electrical conductors when molten or dissolved in water

33

Ionic bonds between ions

must break to melt, boil, etc.

Covalent Bonds

• Solids are usually relatively soft.

– low melting points

– low boiling points

• Properties arise because molecules are not connected to other molecules by bonds, but by intermolecular forces.

• Usually composed of nonmetals

44

Intermolecular

forces, NOTbonds, break to

melt, boil etc.

Quick Review of Bond Types

• Classify the following substances by the type of bond:

• CaF2• CuCl2• NCl3• H2O

• NH4Cl

• K2SO4

55

Types of Bonds

• CaF2 ionic

• CuCl2 ionic

• NCl3 covalent

• H2O covalent

• NH4Cl ionic and covalent

• K2SO4 ionic and covalent

66

Electronegativity

• The ability of an atom in a molecule to attract shared electrons to itself.

• Relative values…no units (4 for F is highest).

• Note the relative trend.

H belongs

between B and C

Non-polar covalent, polar covalent, ionic

• Red = high electron density

• Blue = low electron density

• Green = in between 88

Non-polar covalent

No (or very small

difference in electronegativity

Polar covalent

Electronegativity

difference larger than 0.4-2

Ionic

Electronegativity

difference larger than 2

4 4

0

2.1 4

1.9

1 4

3

Covalent Bond

• There is an optimum distance that will maximize attraction. This is the bond length which is calculated by adding the radii of two atoms.

Bond Lengths and StrengthsWhat is the general trend of bond strength vs. bond length?

What is the trend for number of bonds?

Notice that the trend is not absolute

0

200

400

600

800

1000

1200

100 120 140 160 180 200 220

Bond length (pm)

Bo

nd

Str

en

gth

(K

J/m

ol)

triple

double

single

Bond Lengths and Strengths

In general, triple bonds are stronger (and shorter) than double bonds which are stronger than single bonds.

Also, in general, shorter bonds are stronger than longer bonds. As with anything else in chemistry, there are

exceptions.

Bond Strength

Bond Length

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> >

Lewis Dot Symbols

• Covalent bonding focuses on interactions of valence electrons of two or more atoms

• Can use Lewis Dot Symbols to represent the numbers of valence electrons for each atom (based on electron configurations)

• Usually used for predicting structures of covalent molecules.

Mainly used for covalent

compounds.

Not used for d-block

compounds (which are ionic

anyway).

Electron-Dot Structures

• Used to show how electrons are shared between nonmetals in a covalent bond. Use valence electrons to give each atom an octet, with a few exceptions (e.g., H).

Electron-Dot Structures

• Procedure to give each atom an octet (in most cases):

�Count total valence electrons

• Add or subtract electrons for polyatomic ions

�Draw an atomic skeleton

�Place electron pairs (single bonds) between bonded atoms

�Place remaining electrons on the outside atoms, then the central atom

�Shift electrons, as necessary, to make multiple bonds and satisfy the octet rule and # of valence electrons

1414

A couple of hints for drawing Lewis Structures

• Hydrogen will always be at the end of a molecule– This is because H can only hold 2 electrons in its 1s orbital.

– It will always have a duet.

• Carbon will be a central atom

• Fluorine will always be a peripheral atom

• Most times, carbon will not have lone pairs as a stable molecule.– There are a very few exceptions. CO and CN- are two exceptions.

• Central atoms are usually less electronegative than peripheral atoms

• Only draw multiple bonds if the structure cannot be correctly drawn with single bonds– Always try single bonds first

Exceptions to the Octet Rule

• Odd-Electron Molecules

– Draw a Lewis structure for NO and NO2

– Why does NO2 combine with itself to form N2O4?

– Odd-electron molecules are very reactive. They are called radicals. You will see these much less often than even-number molecules.

• Incomplete Octets

– Certain central molecules don’t need an octet.

– Draw structures for BeCl2, BH3, BF3, AlCl3

1616

Exceptions to the Octet Rule

• What do you do if all of your atoms have an octet…but there are still electrons left over?

• Draw Lewis Structures for the following:

SF4 SF6 IF4+ XeF4

XeF2 PF5 ClF3 BrF5

• Where do we find central atoms that expand their octets?

1717

1818

Cl

The Concept of Resonance

• While Lewis structures do help to predict the structures of many molecules, there are some structures that cannot be satisfactorily represented with a single Lewis structure.

• Lewis formulas don’t always accurately represent bonds. Sometimes it takes multiple formulas to adequately represent the electron distribution.

– Examples: O3, SO3, NO3-, CO3

2-

1919

The Concept of Resonance – Ozone

O OO

Based on the Lewis structure we would expect O – O bond (148 pm) to be longer

than O = O bond (121 pm)

148 pm?121 pm?

Experimental evidence indicates that both bond lengths are exactly

the same, 128 pm.

The Concept of Resonance – Ozone

O OO O OO

The structure of ozone can be best described by using both structures simultaneously.

The structures do not flip “back-and-forth”

Resonance

• How many different valid Lewis formulas can you draw for the following molecules or ions?

SO2 SO3 CO32-

NO3- NCS-

2222

Resonance

2323

Formal Charge

What about structures with non-equivalentresonance structures?

Take the structure for CO2. Are there any resonance structures? If so, how do we choose the correct

structure?

Formal Charge = number of valence electrons in an atom – (number of

nonbonding electrons and ½ of bonding electrons assigned to atom)

Lewis Structure Practice

•CO2

•SO2

•BF3•CH4

•NH4+

•CCl4•NH3

•H2O

•SF2•COCl2

•PCl5•SF4•ClF3•XeF2•OF2•SF6•BrF5•AlCl4

-

•AsH3

2525

• Valence-Shell Electron-Pair Repulsion

• Electron pairs (or groups of pairs) try to avoid one another because of repulsions between like-charged particles

• Regions where electrons are likely to be found are called electron domains:– Lone electron pairs– Single, double, and triple bonds

• Electron domains occur as far apart as possible

• Notice the single, double, and triple bonds each count as ONE electron domain.

2626

Molecular Shapes: The VSEPR Model

• Use A, B, E notation: A = central atom; B = # outer atoms; E = # lone e- pairs

– CH4 = AB4, NH3 = AB3E, H2O = AB2E2

• Can predict the angles between electron domains (charge clouds, areas of electron density):

• 2 domains - linear (180o)

• 3 domains - trigonal planar (120o)

• 4 domains - tetrahedral (109.5o)

• 5 domains - trigonal bipyramidal (90o & 120o)

• 6 domains - octahedral (90o)

2727

VSEPR Theory Five Fundamental VSEPR Geometries

2828

Using these geometries, we can determine two different types of shapes.

Electronic Geometry Molecular Geometry

Shape that is made electron density make from a central

atom.

Shape that is made when we are concerned only with the shape of the bonding

electrons.

O HH

Four areas of electron

density =

Tetrahedral

Angle of bonds

Bent

O HH

Molecular and Electronic Geometries• A = central atom; B = # outer atoms; E = # lone e- pairs

AB2Linear

AB3Trigonal Planar

AB2E

Bent

AB4Tetrahedral

AB3ETrigonal Pyramidal

AB2E2Bent

AB5Trig. Bipyramidal

AB4ESee-Saw

AB3E2T-Shaped

AB6Octahedral

AB5E

Square Pyramid

AB4E2Square Planar

You need to know!

Electronic Geometries

Molecular Geometries

1 lone pair 2 lone pair

What are the molecular geometries of the following

molecules?

•CO2

•SO2

•BF3•CH4

•NH4+

•CCl4•NH3

•H2O

•SF2•COCl2

•PCl5•SF4•ClF3•XeF2•OF2•SF6•BrF5•AlCl4

-

•AsH3

3131

• Should CO2 and SO2 have the same geometry?

3232

Molecular Shapes

• Lewis structures and VSEPR give information about the shapes of molecules and the distributions of electrons. They don’t explain why a bond forms.

• Valence-bond theory considers both bond formation and molecular shape

• Looks at how electrons are shared in a covalent bond

• VB theory considers the atomic orbitals occupied by the valence electrons

3333

Valence Bond Theory Creating Covalent Bonds

We can combine two s orbitals (H2)

Half-filled orbitals overlap so that 2 electrons can share space and form a

covalent bond.

How exactly do orbitals interact to create covalent bonds?

Creating Covalent Bonds

We can combine two porbitals (F2)

We can combine an sand a p orbital (HCl)

Hybridization

zebra + donkey lion + tiger

Zonkey Liger

Hybrid orbitalsAlthough we know that orbitals (s, p, d, f) overlap to form

covalent bonds, there are situations where this simplistic model

seems to fall apart.

Take methane, CH4, for example….carbon has 4 valence electrons to

pair up with 1 valence electron of 4 hydrogens.

2s 2p 2s 2p

The four bonds formed with hydrogen in methane would appear to form from two different types of orbitals.

But we know from experimental evidence that all 4 bonds are exactly the

same! Same length, same strength, and

same bonding angles!

Hybridization in Methane

• Need 4 equivalent orbitals to form the 4 single (σ) bonds (based on VSEPR and experiments)

• Ground state configuration:

E

2s

2p

E

2s

2pPromotion of

an electron

Hybridization in Methane

E sp3

This new model can account for the experimental evidence for methane that suggests that all of the

bonds in CH4 are equivalent.

Hybridization

Each orbital allows a single bond to form

Hybridization in Boron Trifluoride

• Boron trifluoride: BF3• Lewis Dot Structure?

• Boron is central atom

• Electron configuration of boron?

– [He] 2s2 2p1

• VSEPR shape?

– Trigonal planar

Hybridization in Boron Trifluoride

• Need 3 equivalent orbitals to form the 4 single (s) bonds (based on VSEPR and experiments)

• Ground state configuration:

E

2s

2p

E

sp2

pPromotion & hybridization

Notice one unhybridized porbital is left over

Hybridization in Beryllium Chloride

• Beryllium chloride: BeCl2• Lewis Dot Structure?

• Beryllium is central atom

• Electron configuration of beryllium?

– [He] 2s2

• VSEPR shape?

– Linear

Hybridization in Beryllium Chloride

• Need 2 equivalent orbitals to form the 4 single (s) bonds (based on VSEPR and experiments)

• Ground state configuration:

E

2s

2p

E

sp

pPromotion & hybridization

Hybridization in Phosphorus Pentachloride

• Phosphorus pentachloride: PCl5• Lewis Dot Structure?

• Phosphorus is central atom

• Electron configuration of phosphorus?

– [Ne] 3s2 3p3

• VSEPR shape?

– Trigonal bipyramidal

Hybridization in Phosphorus Pentachloride

• How many single bonds in PCl5?

• Where does the 5th orbital come from?

– Expanded octet

– sp3d

E

2s

2p

3d

Promotion & hybridization

sp3dE

3d

Hybridization in Sulfur Hexafluoride

• Sulfur Hexafluoride: SF6• Lewis Dot Structure?

• Sulfur is central atom

• Electron configuration of sulfur?

– [Ne] 3s2 3p4

• VSEPR shape?

– Octahedral

Hybridization in Sulfur Hexafluoride

• Need 6 equivalent orbitals to form the 6 single bonds.

• Where does the 6th orbital come from?

– Expanded octet

– sp3d2

2s

2p

3d

Promotion & hybridization

sp3d2

E

3d

E

Hybridization and Geometry

Linear

Trigonal Planar

Tetrahedral

Trigonal Bipyramidal

Octahedral

sp

sp2

sp3

sp3d

sp3d2

Once the electronic geometry of a molecule is known, the hybridization

can be predicted

Hybridization in Multiple Bonds: Ethylene

• Ethylene: CH2CH2

• Lewis Dot Structure?

• Carbon is central atom

• Electron configuration of carbon?

– [He] 2s2 2p2

• VSEPR shape?

– Trigonal planar

Hybridization in Ethylene

E

2s

2pPromotion &

hybridization Esp2

Based on the Lewis structure, there should be a double bond.

How does a double bond form?

2p

Each CarbonEach Carbon

Hybridization in Ethylene

sp2 sp2 2p2p

sigma bond (σσσσ): direct overlap of bonding orbitals.

pi bond (ππππ): “sideways” overlap of bonding orbitals

(usually between unhybridized p orbitals)

bond formed by p-poverlap

Hybridization in EthyleneUnhybridized p orbital from C

3 Hybridized sp2

orbital from C

s orbital from H

Hybridization in Acetylene

• Acetylene: HCCH

• Lewis Dot Structure?

• Carbon is central atom

• Electron configuration of carbon?

– [He] 2s2 2p2

• VSEPR shape?

– Linear

sp sp 2p 2p

bond formed by p-p overlap (ππππ)

bond formed by p-p overlap (ππππ)

Hybridization in Acetylene

Hybridization in Acetylene

• sp hybridization accounts for single (σ) bonds, but what about the unhybridized 2p orbitals?

• Those 2 unhybridized 2p orbitals help make the two π bonds.

2 Unhybridized p orbitals from C

s orbital from H

2 Hybridized sp orbitals from C

Summary

• According to Valence Bond Theory, covalent bonds form when:

– there are two electrons in each orbital; one electron from each atom

– these two orbitals “overlap”

• The number of hybridized orbitals equals the number of atomic orbitals that are combined.

– sp � 2 orbitals combined (BeCl2 and HCN)

– sp2 � 3 orbitals combined (BF3 and CH2O)

– sp3 � 4 orbitals combined (CH4)

– sp3d � 5 orbitals combined (PCl5)

– sp3d2 � 6 orbitals combined (SF6)

Summary

• The number of hybridized orbitals equals the number of electron domains around a central atom (starting with s)– sp, sp2, sp3, sp3d, sp3d2

• A single bond has 1 σ bond (the same goes for a lone pair of electrons)

• A double bond has 1 σ bond and 1 π bond

• A triple bond has 1 σ bond and 2 π bonds

• Unhybridized p orbitals participate in π bonding (to make double and triple bonds)