1 Chapter 8 Covalent bonding. 2 Covalent Bonding A metal and a nonmetal transfer electrons –An...

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Chapter 8

Covalent bonding

2

Covalent BondingA metal and a nonmetal transfer

electrons– An ionic bond

Two metals just mix and don’t react– An alloy

What do two nonmetals do?– Neither one will give away an electron– So they share their valence electrons– This is a covalent bond

3

Covalent bondingMakes molecules

– Specific atoms joined by sharing electrons

Two kinds of molecules:Molecular compound

– Sharing by different elementsDiatomic molecules

– Two of the same atom– O2 N2

4

Diatomic elementsThere are 8 elements that always form

molecules

H2 , N2 , O2 , F2 , Cl2 , Br2 , I2 , and At2

Oxygen by itself means O2

The –ogens and the –ines

1 + 7 pattern on the periodic table

5

1 and 7

6

Molecular compoundsTend to have low melting and boiling

points

Have a molecular formula which shows type and number of atoms in a molecule

Not necessarily the lowest ratio

C6H12O6

Formula doesn’t tell you about how atoms are arranged

7

Polar BondsWhen the atoms in a bond are the

same, the electrons are shared equally.

This is a nonpolar covalent bond.

When two different atoms are connected, the electrons may not be shared equally.

This is a polar covalent bond.

How do we measure how strong the atoms pull on electrons?

8

ElectronegativityA measure of how strongly the atoms

attract electrons in a bond.The bigger the electronegativity

difference the more polar the bond.Use table 12-3 Pg. 2850.0 - 0.4 Covalent nonpolar0.5 - 1.0 Covalent moderately polar1.0 -2.0 Covalent polar>2.0 Ionic

9Chapter 9: Chemical Bonds 9

Electronegativity

Electronegativity (EN) is a measure of the ability of an atom to attract bonding electrons to itself

EOS

The greater the electronegativity of an atom in a molecule, the more strongly it attracts the electrons in a covalent bond

10Chapter 9: Chemical Bonds 10

Pauling’s Electronegativities

EOS

Electronegativity Illustrated

11Chapter 9: Chemical Bonds 11

Electronegativity Differenceand Bond Type

Two identical atoms have the same electronegativity and share a bonding electron pair equally. This is called a nonpolar covalent bond

Example: chlorine gas

EOS

All homonuclear diatomic molecules have nonpolar covalent bonds:

H2, N2, O2, F2, Cl2, Br2, I2

12Chapter 9: Chemical Bonds 12

Electronegativity Differenceand Bond Type

In covalent bonds between atoms with somewhat larger electronegativity differences, electron pairs are shared unequally. This is called a polar covalent bond

Example: hydrogen chloride gas, HCl

EOS

The electrons are drawn closer to the atom of higher electronegativity, Cl

13

Covalent BondingElectrons are shared by atoms.

These are two extremes.

In between are polar covalent bonds.

The electrons are not shared evenly.

One end is slightly positive, the other negative.

Indicated using small delta

14

How to show a bond is polar Isn’t a whole charge just a partial charge

means a partially positive

means a partially negative

The Cl pulls harder on the electrons

The electrons spend more time near the Cl

H Cl

15

H - F+ -

H - F

+-H - F+

-

H - F

+-

H - F +-

H - F+-

H - F

+-

H - F

+-

16 © 2009, Prentice-Hall, Inc.

Polar Covalent Bonds

When two atoms share electrons unequally, a bond dipole results.

The dipole moment, , produced by two equal but opposite charges separated by a distance, r, is calculated:

= Qr

It is measured in debyes (D).

17Chapter 9: Chemical Bonds 17

Electronegativity Differenceand Bond Type

With still larger differences in electronegativity, electrons may be completely transferred from metal to nonmetal atoms to form ionic bonds

Example: sodium chloride, NaCl

EOS

18Chapter 9: Chemical Bonds 18

Electronegativity Differences

EOS

19Chapter 9: Chemical Bonds 19

Electron Distributions and Covalent Bonds

Symmetric distribution

EOS

AsymmetricDistribution

20

Electronegativity difference

Bond Type

Zero

Intermediate

Large

Covalent

Polar Covalent

Ionic

Co

valent C

haracter

decreases

Ion

ic Ch

aracter increases

21

How does H2 form?The nuclei repel

++

22

How does H2 form?

++

The nuclei repel

But they are attracted to electrons

They share the electrons

23Chapter 9: Chemical Bonds 23

The Lewis Theory ofChemical Bonding: An

OverviewElectrons, particularly valence electrons, play a fundamental role in chemical bonding

EOS

In losing, gaining, or sharing electrons to form chemical bonds, atoms tend to acquire the electron configurations of noble gases

24Chapter 9: Chemical Bonds 24

Lewis Symbols

Valence electrons are shown by dots around the element symbol

EOS

Use rules of electron configurations when forming dot structures …e.g., electrons remain unpaired if possible

25

Covalent bondsNonmetals hold onto their valence

electrons.

They can’t give away electrons to bond.

Still need noble gas configuration.

Get it by sharing valence electrons with each other.

By sharing both atoms get to count the electrons toward noble gas configuration.

26

Covalent bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons Both end with full orbitals

F F

27

Covalent bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons Both end with full orbitals

F F8 Valence electrons

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Covalent bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons Both end with full orbitals

F F8 Valence electrons

29

Lewis StructureShows how the valence electrons are

arranged.One dot for each valence electron.A stable compound has all its atoms with

a noble gas configuration.Hydrogen follows the duet rule.The rest follow the octet rule.Bonding pair is the one between the

symbols.

30Chapter 9: Chemical Bonds 30

Multiple Covalent BondsThe covalent bond in which one pair of electrons is shared is called a single bond

e.g., H : Cl or H—Cl

Double bonds have two shared pairs of electrons

EOS

Triple bonds have three shared pairs of electrons

31

RulesSum the valence electrons.

Use a pair to form a bond between each pair of atoms.

Arrange the rest to fulfill the octet rule (except for H and the duet).

H2O

A line can be used instead of a pair.

32

A useful equation(happy-have) / 2 = bonds

CO2 C is central atom

POCl3 P is central atom

SO42-

S is central atom

SO32-

S is central atom

PO43-

P is central atom

SCl2 S is central atom

33

Exceptions to the octetBH3

Be and B often do not achieve octetHave less than an octet, for electron

deficient molecules.

SF6

Third row and larger elements can exceed the octet

Use 3d orbitals?

I3-

34

Exceptions to the octetWhen we must exceed the octet, extra

electrons go on central atom.

(Happy – have)/2 won’t work

ClF3

XeO3

ICl4-

BeCl2

35Chapter 9: Chemical Bonds 35

Writing Lewis Structures

Hydrogen atoms are terminal atoms (bonded to only one other atom)

EOS

The central atom of a structure usually has the lowest electronegativity and the terminal atoms (except H) generally have higher electronegativities

36

Single Covalent BondA sharing of two valence electrons.

Only nonmetals and Hydrogen.

Different from an ionic bond because they actually form molecules.

Two specific atoms are joined.

In an ionic solid you can’t tell which atom the electrons moved from or to.

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How to show how they formed It’s like a jigsaw puzzle.

I have to tell you what the final formula is.

You put the pieces together to end up with the right formula.

For example- show how water is formed with covalent bonds.

38

Water

H

O

Each hydrogen has 1 valence electron

and wants 1 more

The oxygen has 6 valence electrons

and wants 2 more

They share to make each other “happy”

39

WaterPut the pieces together

The first hydrogen is happy

The oxygen still wants one more

H O

40

WaterThe second hydrogen attaches

Every atom has full energy levels

H OH

41

Multiple BondsSometimes atoms share more than one

pair of valence electrons.

A double bond is when atoms share two pair (4) of electrons.

A triple bond is when atoms share three pair (6) of electrons.

42

Carbon dioxideCO2 - Carbon is central

atom ( I have to tell you)

Carbon has 4 valence electrons

Wants 4 more

Oxygen has 6 valence electrons

Wants 2 moreO

C

43

Carbon dioxideAttaching 1 oxygen leaves the oxygen 1

short and the carbon 3 short

OC

44

Carbon dioxide Attaching the second oxygen leaves

both oxygen 1 short and the carbon 2 short

OCO

45

Carbon dioxide The only solution is to share more

OCO

46

Carbon dioxide The only solution is to share more

OCO

47

Carbon dioxide The only solution is to share more

OCO

48

Carbon dioxide The only solution is to share more

OCO

49

Carbon dioxide The only solution is to share more

OCO

50

Carbon dioxide The only solution is to share more

OCO

51

Carbon dioxide The only solution is to share more Requires two double bonds Each atom gets to count all the atoms in the

bond

OCO

52

Carbon dioxide The only solution is to share more Requires two double bonds Each atom gets to count all the atoms in the

bond

OCO8 valence electrons

53

Carbon dioxide The only solution is to share more Requires two double bonds Each atom gets to count all the atoms in the

bond

OCO8 valence electrons

54

Carbon dioxide The only solution is to share more Requires two double bonds Each atom gets to count all the atoms in the

bond

OCO

8 valence electrons

55

How to draw themTo figure out if you need multiple bondsAdd up all the valence electrons.Count up the total number of electrons to

make all atoms happy.Subtract.Divide by 2Tells you how many bonds - draw them.Fill in the rest of the valence electrons to

fill atoms up.

56

ExamplesNH3

N - has 5 valence electrons wants 8

H - has 1 valence electrons wants 2

NH3 has 5+3(1) = 8

NH3 wants 8+3(2) = 14

(14-8)/2= 3 bonds

4 atoms with 3 bonds

N

H

57

N HHH

ExamplesDraw in the bonds

All 8 electrons are accounted for

Everything is full

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ExamplesHCN C is central atom

N - has 5 valence electrons wants 8

C - has 4 valence electrons wants 8

H - has 1 valence electrons wants 2

HCN has 5+4+1 = 10

HCN wants 8+8+2 = 18

(18-10)/2= 4 bonds

3 atoms with 4 bonds -will require multiple bonds - not to H

59

HCNPut in single bonds

Need 2 more bonds

Must go between C and N

NH C

60

HCN Put in single bonds Need 2 more bonds Must go between C and N Uses 8 electrons - 2 more to add

NH C

61

HCN Put in single bonds Need 2 more bonds Must go between C and N Uses 8 electrons - 2 more to add Must go on N to fill octet

NH C

62

Where do bonds go?Think of how many electrons they are

away from noble gas.

H should form 1 bond- always

O should form 2 bonds – if possible

N should form 3 bonds – if possible

C should form 4 bonds– if possible

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PracticeDraw electron dot diagrams for the

following.

PCl3

H2O2

CH2O

C3H6

64

Another way of indicating bonds

Often use a line to indicate a bond

Called a structural formula

Each line is 2 valence electrons

H HO =H HO

65

Structural Examples

H C N

C OH

H

C has 8 electrons because each line is 2 electrons

Ditto for N

Ditto for C here

Ditto for O

66

Coordinate Covalent BondWhen one atom donates both electrons

in a covalent bond.

Carbon monoxide

CO

OC

67

Coordinate Covalent Bond When one atom donates both electrons

in a covalent bond. Carbon monoxide CO

OC

68

Coordinate Covalent Bond When one atom donates both electrons

in a covalent bond. Carbon monoxide CO

OCOC

69

How do we know ifHave to draw the diagram and see what

happens.

Often happens with polyatomic ions

If an element has the wrong number of bonds

70

Polyatomic ionsGroups of atoms held by covalent

bonds, with a charge

Can’t build directly, use (happy-have)/2

Have number will be different

Surround with [ ], and write charge

NH42+

ClO21-

71

ResonanceWhen more than one dot diagram with

the same connections is possible.Choice for double bond

NO2-

Which one is it?Does it go back and forth?Double bonds are shorter than single

In NO2- all the bonds are the same

length

72

Formal ChargeFor molecules and polyatomic ions that

exceed the octet there are several different structures.

Use charges on atoms to help decide which.

Trying to use the oxidation numbers to put charges on atoms in molecules doesn’t work.

73

Formal ChargeThe difference between the number of valence electrons on the free atom and that assigned in the molecule or ion.

We count half the electrons in each bond as “belonging” to the atom.

SO4-2

Molecules try to achieve as low a formal charge as possible.

Negative formal charges should be on electronegative elements.

74Chapter 9: Chemical Bonds 74

Formal Charge

EOS

Formal charge is the difference between the number of valence electrons in a free (uncombined) atom and the number of electrons assigned to that atom when bonded to other atoms in a Lewis structure

75Chapter 9: Chemical Bonds 75

Formal Charge

Usually, the most plausible Lewis structure is one with no formal charges

When formal charges are required, they should be as small as possible and negative formal charges should appear on the most electronegative atoms

EOS

Adjacent atoms in a structure should not carry formal charges of the same sign

76

ExamplesXeO3

NO43-

SO2Cl2

77

Resonance It is a mixture of both, like a mule.

CO32-

78

Bond Dissociation EnergyThe energy required to break a bond

C - H + 393 kJ C + H

Double bonds have larger bond dissociation energies than single

Triple even larger

– C-C 347 kJ

– C=C 657 kJ

– C≡C 908 kJ

79

Bond Dissociation EnergyThe larger the bond energy, the harder

it is to break

Large bond energies make chemicals less reactive.