Chapter 7. Comparing Ionic and Covalent Bonds. Ionic Bonds. Types of Bonds. Quick Review of Bond Types. Covalent Bonds

Chapter 7 Chapter 7 Chapter 7 Chapter 7

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Covalent Bonds Covalent Bonds and Molecular and Molecular Structure Structure

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

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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

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Intermolecular forces, NOT bonds, break to

melt, boil etc.

Quick Review of Bond Types

• Classify the following substances by the type of bond:

• CaF

• CuCl

• NCl

• H

O

• NH

Cl

• K

SO

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Types of Bonds

• CaF

ionic

• CuCl

ionic

• NCl

covalent

• H

O covalent

• NH

Cl ionic and covalent

• K

SO

ionic and covalent

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Electronegativity

• The ability of an atom in a molecule to attract sharedelectrons 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

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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 Strengths

What 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 le ngth (pm )

Bond Strength (KJ/mol)

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

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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 NO₂combine 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

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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:

SF

SF

IF

XeF

XeF

PF

ClF

BrF

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

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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-, CO32-

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The Concept of Resonance – Ozone

O O O

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 O

O O O O

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?

SO

SO

CO

NO

NCS

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Resonance

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Formal Charge

What about structures with non-equivalent resonance 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

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• 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.

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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 (180^o)

• 3 domains - trigonal planar (120^o)

• 4 domains - tetrahedral (109.5^o)

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

• 6 domains - octahedral (90^o)

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VSEPR Theory Five Fundamental VSEPR Geometries

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Using these geometries, we can determine two different typesof 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 H H

Four areas of electron density = Tetrahedral

Angle of bonds Bent

O H H

Molecular and Electronic Geometries

• A = central atom; B = # outer atoms; E = # lone e^-pairs

AB₂ Linear

AB₃ Trigonal Planar

AB₂E Bent AB₄

Tetrahedral

AB₃E Trigonal Pyramidal

AB₂E₂ Bent AB₅

Trig. Bipyramidal

AB₄E See-Saw

AB₃E₂ T-Shaped AB₆

Octahedral

AB₅E Square Pyramid

AB₄E₂ Square Planar You need to know!

Electronic Geometries

Molecular Geometries

1 lone pair 2 lone pair

What are the molecular geometries of the following molecules?

•CO

•SO

•BF

•CH

•NH

•CCl

•NH

•H

O

•SF

•COCl

•PCl

•SF

•ClF

•XeF

•OF

•SF

•BrF

•AlCl

•AsH

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• Should CO

and SO

have the same geometry?

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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

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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 p orbitals (F₂)

We can combine an s and a p orbital (HCl)

Hybridization

zebra + donkey lion + tiger

Zonkey Liger

Hybrid orbitals

Although 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

2p

Promotion 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 CH4are equivalent.

Hybridization

Each orbital allows a single bond to form

Hybridization in Boron Trifluoride

• Boron trifluoride: BF

• Lewis Dot Structure?

• Boron is central atom

• Electron configuration of boron?

– [He] 2s²2p¹

• 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

Promotion & p

hybridization

Notice one unhybridized p orbital is left over

Hybridization in Beryllium Chloride

• Beryllium chloride: BeCl

• 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

Promotion & p

hybridization

Hybridization in Phosphorus Pentachloride

• Phosphorus pentachloride: PCl

• Lewis Dot Structure?

• Phosphorus is central atom

• Electron configuration of phosphorus?

– [Ne] 3s²3p³

• VSEPR shape?

– Trigonal bipyramidal

Hybridization in Phosphorus Pentachloride

• How many single bonds in PCl

?

• Where does the 5

orbital come from?

– Expanded octet – sp³d

E 2s

2p 3d

Promotion &

hybridization

sp³d E

3d

Hybridization in Sulfur Hexafluoride

• Sulfur Hexafluoride: SF

• Lewis Dot Structure?

• Sulfur is central atom

• Electron configuration of sulfur?

– [Ne] 3s²3p⁴

• VSEPR shape?

– Octahedral

Hybridization in Sulfur Hexafluoride

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

• Where does the 6

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: CH

CH

• Lewis Dot Structure?

• Carbon is central atom

• Electron configuration of carbon?

– [He] 2s²2p²

• VSEPR shape?

– Trigonal planar

Hybridization in Ethylene

E 2s

2p Promotion &

hybridization E

sp² Based on the Lewis structure, there should be a double bond.

How does a double bond form?

2p Each Carbon Each Carbon

Hybridization in Ethylene

sp² sp² 2p 2p

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-p overlap

Hybridization in Ethylene

Unhybridized 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] 2s²2p²

• VSEPR shape?

– Linear

sp sp 2p 2p

bond formed by p-p overlap (πππ)π

bond formed by p-p overlap (πππ)π

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 (BeCl2and HCN) – sp² 3 orbitals combined (BF3and 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 porbitals participate in π bonding (to make double and triple bonds)