Chemical Bonding

Chapter 6

Chemistry chapter 6 2

Chemical bond• A mutual electrical attraction between

the nuclei and valence electrons of different atoms that binds them together

Chemistry chapter 6 3

Types of bonds• Ionic

• electrons are transferred from one atom to another, forming ions

• Electrical attraction of large numbers of cations and anions holds the compound together

• Covalent – electron pairs are shared between the atoms

Chemistry chapter 6 4

Ionic Bonds• Ions get very close together and are attracted

to each other.• Example: NaCl

Na+

Cl-

Chemistry chapter 6 5

Covalent Bonds• The atoms share electrons, and their

electron clouds overlap.• Example: carbon monoxide

CO

Chemistry chapter 6 6

Electronegativity difference• Can be used to predict how two atoms

will react with each other.• Has no units, because it is a

comparison.• When it is high, the bond is ionic.• When it is low, the bond is covalent.

Chemistry chapter 6 7

Ionic – covalent continuum• Real bonds are usually sort of

ionic and sort of covalent.• See figure 6-2 on page 162.• Difference of 1.7 or less is

considered covalent• 50% ionic character or less

• Difference of more than 1.7 is considered ionic• More than 50% ionic character

Chemistry chapter 6 8

Nonpolar-covalent bond• Bonding electrons are shared equally• Balanced distribution of electrical

charge• 0 to 5% ionic character• 0 to 0.3 electronegativity difference

Chemistry chapter 6 9

Polar• Having uneven charge distribution

Chemistry chapter 6 10

Polar-covalent bond• The bonded atoms have unequal

attraction for the shared electrons• 5 to 50% ionic character• 0.3 to 1.7 electronegativity difference

Chemistry chapter 6 11

Polar and nonpolar bonds

• Electron density is greater around Cl than H because it has greater electronegativity

• d+ and d- indicate partial charge• Electron cloud drawings

nonpolar polar

-d+d

Chemistry chapter 6 12

Comparing bond types• Comparison between ionic and covalent

bonding

Chemistry chapter 6 13

Discuss• Page 163

• Sample problem• Practice problem• Section review 1 and 2

Chemistry chapter 6 14

Molecule• A neutral group of atoms that are held

together by covalent bonds• A single molecule is capable of existing

on its own• May consist of atoms of the same

element or different elements

Chemistry chapter 6 15

Molecular compound• Compound whose simplest units are

molecules

Chemistry chapter 6 16

Chemical formula• Uses atomic symbols and subscripts to

indicate the relative numbers of atoms of each kind in a chemical compound.

Chemistry chapter 6 17

Molecular formula• Shows the types and numbers of atoms

in a single molecule of a molecular compound

• H2O

• O2

• C12H22O11

Chemistry chapter 6 18

Diatomic molecule• Contains only two atoms

Chemistry chapter 6 19

Bonding• Bonding makes atoms be at lower

potential energy levels, which is favored by nature.

• Attractive forces (nucleus-electron) are balanced by repulsive forces (nucleus-nucleus and electron-electron).

Chemistry chapter 6 20

Bond length• Average distance between bonded

nuclei at their lowest potential energy

Chemistry chapter 6 21

Bond energy• The amount of energy that must be

added to break a chemical bond and form neutral isolated atoms.

• The same amount of energy that the atoms had to release when they bonded.

Chemistry chapter 6 22

Octet rule• Chemical compounds tend to form so

that each atom, by gaining, losing or sharing electrons, has an octet of electrons in its highest occupied energy level

• The outermost s and p sublevels want to be full

Chemistry chapter 6 23

Exceptions to octet rule• Hydrogen only needs two electrons, not

eight• Boron tends to only be surrounded by

six electrons.• Some atoms can have more than eight

when combining with fluorine, oxygen, and chlorine. This expanded valence also includes the d orbitals.

Chemistry chapter 6 24

Electron dot diagrams• Usually only the valence electrons are

involved in chemical reactions.• Electron dot diagrams allow us to draw

these electrons around the symbol for an element.

Chemistry chapter 6 25

Drawing electron dot diagrams

1. The element’s symbol represents the nucleus and all the electrons in the inner energy levels.

2. Write the noble gas configuration of the element. Add the superscripts of the s and p orbitals to get the number of valence electrons.

Chemistry chapter 6 26

3. Draw dots on the sides to represent electrons.

Chemistry chapter 6 27

Examples• Examples

Chemistry chapter 6 28

Molecules• Can be represented by dot diagrams• Example: HCl• Shared pairs can be replaced by long

dashes

Chemistry chapter 6 29

Lone pair• Unshared pair of electrons• Not involved in bonding

Chemistry chapter 6 30

Discuss• Read sample problem 6-2 on page 170

Chemistry chapter 6 31

Lewis diagram• Formulas in which atomic symbols

represent the nuclei and inner-shell electrons• Dot diagrams with dots or dashes for

shared electrons

Chemistry chapter 6 32

Structural diagrams• Leaves out the unshared pairs• Indicates the kind, number,

arrangement, and bonds of atoms

Chemistry chapter 6 33

Single bond• Covalent bond produced when one pair

of electrons is shared

Chemistry chapter 6 34

Discuss• Sample problem 6-3 on pages 171-172

Chemistry chapter 6 35

Double bond• Covalent bond produced when two

pairs of electrons are shared• Example: CO2

Chemistry chapter 6 36

Triple bond• Covalent bond formed when three pairs

of electrons are shared• Example: N2

Chemistry chapter 6 37

Multiple bonds• Double or triple bonds• Have higher bond energies• Have shorter bond lengths

Chemistry chapter 6 38

Resonance structures• Bonding in molecules or ions that

cannot be correctly represented by a Lewis structure

• A double-headed arrow is placed between the two possibilities

• Example: ozone

Chemistry chapter 6 39

Discuss• Covalent Bonds• More Covalent Bonds• Sample problem 6-4 on page 174

Chemistry chapter 6 40

Examples• Draw a Lewis Structure and the structural

formula for each of the following: • H2O

• CH4

• CH4O

• O2

• C2H2

• C2H4

Chemistry chapter 6 41

Discuss• Section Review on page 175

Chemistry chapter 6 42

Ionic compound• Composed of positive and negative ions that

are combined so that the numbers of positive and negative charges are equal

• Usually crystalline solids• Not composed of independent units like

molecules• Ionic bonding• More ionic bonding

Chemistry chapter 6 43

Chemical formula• For ionic compounds

• Shows the simplest ratio of the ions that gives no net charge

Chemistry chapter 6 44

Formula unit• The simplest collection of atoms from

which an ionic compound’s formula can be established.

• For sodium chloride, NaCl• The simplest way to show a one to one

ratio

Chemistry chapter 6 45

Ratio of ions• Depends on the charges of the ions

combined.• Na+ and Cl- form NaCl• Mg2+ and F- form MgF2

Chemistry chapter 6 46

Forming ionic compounds• Examples:• NaCl• MgF2

Chemistry chapter 6 47

Crystal lattice• Ions minimize their potential energy by

forming an orderly arrangement

Chemistry chapter 6 48

Lattice energy• The energy released when one mole of

an ionic crystalline compound is formed from gaseous ions.

• Used to compare bond strengths

Chemistry chapter 6 49

Molecular vs. ionic compounds

• Forces between molecules are weaker than ionic bonding forces

• Molecular compounds have lower melting points

• Ionic compounds are hard but brittle.• Ionic compounds conduct electricity

when melted or dissolved in solution

Chemistry chapter 6 50

Polyatomic ions• A charged group of covalently bonded

atoms.• Form ionic bonds just like other ions.• Examples

• Ammonium ion NH4+

• Sulfate ion SO42-

Chemistry chapter 6 51

Discuss• Use electron-dot notation to

demonstrate the formation of ionic compounds involving• Li and Cl• Ca and I

• What basic unit are molecular compounds composed of? Ionic compounds?

Chemistry chapter 6 52

Metals• High electrical conductivity.• Metals form crystals and their orbitals

overlap.• All the atoms have the same attraction

for their electrons, so the electrons can move easily from one atom to another.

Chemistry chapter 6 53

Valence electrons• Metals have very few• Their p sublevels are empty• Transition metals also have many

vacant d sublevels

Chemistry chapter 6 54

Metallic bonding• Delocalized electrons – don’t belong to

one specific atom• Sea of electrons formed around the

metal atoms• Metallic bonding results from the

attraction between metal atoms and the surrounding sea of electrons

Chemistry chapter 6 55

Metallic properties• High electrical and thermal conductivity• Ability to absorb a wide range of light

frequencies• Electrons enter excited states and then

return to ground states• This makes metals look shiny

Chemistry chapter 6 56

Metallic properties• Metals bond the same way in all

directions• Their layers can slide past each other

without breaking• Malleabiltiy - the ability to be hammered or

beaten into thin sheets• Ductility – the ability to be drawn, pulled or

extruded into a thin wire

Chemistry chapter 6 57

Metallic bond strength• When there is a bigger nucleus and

more electrons, the bond is stronger.• Stronger bonds have higher heats of

vaporization

Chemistry chapter 6 58

Discuss• Describe the electron-sea model of

metallic bonding.• What is the relationship between

metallic bond strength and heat of vaporization?

• Explain why most metals are malleable and ductile but ionic crystals are not.

Chemistry chapter 6 59

Molecular polarity• Uneven distribution of molecular charge• Determined by the polarity of each bond

and the geometry of the molecule

Chemistry chapter 6 60

Molecular geometry• There are two equally successful

theories• One accounts for molecular bond

angles• The other describes the orbitals that

contain the valence electrons

Chemistry chapter 6 61

VSEPR theory• Valence-shell, electron-pair repulsion• The repulsion between the sets of

valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible

• To predict geometry, one must consider the location of all electron pairs around the bonded atoms.

Chemistry chapter 6 62

Linear molecules• A is the central atom• B is an atom bonded to it• Formulas like AB2 form linear

molecules, with the B atoms on opposite sides of the A atom

Chemistry chapter 6 63

Trigonal Planar

• AB3

• A is in the center.• The 3 Bs make a triangle around it.• The molecule is flat (planar)

Chemistry chapter 6 64

Tetrahedral

• AB4

• A pyramid with A at the center and a B at each corner

Chemistry chapter 6 65

When there are lone pairs• The repel the other electrons just like

shared pairs do• Use an E to represent the unshared

pairs

Chemistry chapter 6 66

Bent or angular

• AB2E

• Like trigonal planar, but without the top of the triangle

Chemistry chapter 6 67

Trigonal pyramidal• AB3E

• Like tetrahedral, but without the top of the pyramid

• A is at the center, the Bs are the base, and E is at the top

Chemistry chapter 6 68

Bent or angular

• AB2E2

• Like a tetrahedral, but only the atoms (not the unshared pairs) determine the shape

Chemistry chapter 6 69

Trigonal bipyramidal

• AB5

• Like two pyramids on top of each other

Chemistry chapter 6 70

Octahedral

• AB6

• 8 sided shape

Chemistry chapter 6 71

VSEPR• Double and triple bonds are treated the

same way as single bonds.• Use Lewis structures and table 6-5 on

page 186 to predict molecule shapes

Chemistry chapter 6 72

Hybridization• Explains how orbitals get rearranged

when atoms form covalent bonds• Describes how atoms are bonded

Chemistry chapter 6 73

Hybrid orbitals• Orbitals of equal energy produced by

the combining of two or more orbitals on the same atom

• The number of hybrid orbitals equals the number of orbitals they were made from

Chemistry chapter 6 74

Example: Methane

• CH4

• To get 4 equal orbitals, the s and p orbitals combine.

• They form 4 sp3 orbitals

Chemistry chapter 6 75

Other hybrid orbitals

Chemistry chapter 6 76

Intermolecular forces• The forces of attraction between

molecules• Usually weaker than covalent, ionic,

and metallic bonds• Can be measured by boiling point

• Higher boiling point means stronger intermolecular forces

Chemistry chapter 6 77

Dipole• Equal but opposite charges separated

by a short distance• Direction is from the positive pole to the

negative pole

ClH

Chemistry chapter 6 78

Dipole-dipole forces• Forces of attraction between polar

molecules• Strongest intermolecular forces• Only act over short distances

Chemistry chapter 6 79

More than two atoms• Polarity of the molecule depends on the

polarity and orientation of each bond

Chemistry chapter 6 80

Chemistry chapter 6 81

Induced dipoles• A polar molecule can induce a dipole in

a nonpolar molecule by attracting its electrons.

• It is temporary

Chemistry chapter 6 82

Hydrogen bonding• A hydrogen atom in a compound with a

very electronegative atom is attracted to an unshared pair of electrons in a nearby molecule.

• Examples:• Water (H2O)• Hydrogen fluoride (HF)• Ammonia (NH3)

Chemistry chapter 6 83

Hydrogen bonding• Causes higher boiling points

Chemistry chapter 6 84

Polarity and Hydrogen Bonding

• Polarity and Hydrogen Bonding

Chemistry chapter 6 85

London dispersion forces• The intermolecular attractions caused

by the constant motion of electrons and the creation of instantaneous dipoles

• Act between all atoms and molecules• The only intermolecular forces among

noble gas atoms and nonpolar molecules• Low boiling points

Chemistry chapter 6 86

discuss• What is the difference between

intramolecular forces and intermolecular forces?

• What is a dipole? In what kind of molecule are dipoles common?

• What is hydrogen bonding?• What are London dispersion forces?• Molecular Geometry review