Chemical Bonding (Predicting Bond Types) Lewis (Electron...

Chemical Bonding (Predicting Bond Types)

Lewis (Electron) Dot Diagrams

Binary Molecular Nomenclature

Exceptions to the Octet Rule

Coordinate Covalent Bonding

Resonance Structures

Molecular Shapes and Polarity

Intermolecular Forces of Attraction

What is a chemical bond?

A chemical bond is a strong

attractive force between

atoms or ions in a chemical

compound.

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Why do elements form chemical bonds?

1. Uncombined elements have relatively

high potential energy.

2. Atoms will gain, lose or share valence

electrons in order to chemically combine

with other atoms.

3. By combining with other atoms, atoms

decrease potential energy and create

more stable arrangements.

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What two factors determine whether or not a chemical bond will form?

1. the electron configurations of

the atoms involved

2. the attraction the atoms have

for electrons

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How is the type of chemical bond formed between two atoms determined?

The type of chemical bond formed

depends upon the degree to which

the valence electrons are shared

between the atoms.

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Covalent Bonding

In a covalent bond, valence

electrons are shared by the atoms.

Covalent bonds can be nonpolar or

polar.

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Nonpolar vs. Polar

In a nonpolar covalent bond, electrons

are shared equally. Bonding which

occurs between two atoms of the same

element is an example of nonpolar

covalent bonding.

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

In a polar covalent bond, electrons are

shared unequally.

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Ionic Bonding

In an ionic bond, valence electrons

are transferred between atoms.

One atom gains electrons to form a

negative ion (anion) and the other

atom loses electrons to form a

positive ion (cation).

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Ionic Bonding Which category of elements tends to

gain electrons and form negative ions

(anions)?

nonmetals

Which category of elements tends to

lose electrons and form positive ions

(cations)?

metals

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Using differences in electronegativity to determine bond type

Electronegativity is a measure of an atom’s ability

to attract electrons when chemically combining

with another element.

The higher the electronegativity value, the stronger

the attraction the atom has for another atom’s

electrons.

The degree to which bonding between atoms of

two elements is ionic or covalent can be estimated

by calculating the difference in the elements’

electronegativities (ΔEN).

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Using differences in electronegativity to determine bond type

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Type of

Bond

∆EN (Difference in

Electronegativity)

Nonpolar

Covalent ≤0.2

Polar

Covalent 0.2 to 1.7

Ionic ≥1.7

1

H

2.1

Periodic Table of Electronegativities 2

He

-

3

Li

1.0

4

Be

1.5

5

B

2.0

6

C

2.5

7

N

3.0

8

O

3.5

9

F

4.0

10

Ne

-

11

Na

0.9

12

Mg

1.2

13

Al

1.5

14

Si

1.8

15

P

2.1

16

S

2.5

17

Cl

3.0

18

Ar

-

19

K

0.8

20

Ca

1.0

21

Sc

1.3

22

Ti

1.5

23

V

1.6

24

Cr

1.6

25

Mn

1.5

26

Fe

1.8

27

Co

1.9

28

Ni

1.8

29

Cu

1.9

30

Zn

1.6

31

Ga

1.6

32

Ge

1.8

33

As

2.0

34

Se

2.4

35

Br

2.8

36

Kr

3.0

37

Rb

0.8

38

Sr

1.0

39

Y

1.2

40

Zr

1.4

41

Nb

1.6

42

Mo

1.8

43

Tc

1.9

44

Ru

2.2

45

Rh

2.2

46

Pd

2.2

47

Ag

1.9

48

Cd

1.7

49

In

1.7

50

Sn

1.8

51

Sb

1.9

52

Te

2.1

53

I

2.5

54

Xe

2.6

55

Cs

0.7

56

Ba

0.9

57

La

1.1

72

Hf

1.3

73

Ta

1.4

74

W

1.7

75

Re

1.9

76

Os

2.2

77

Ir

2.2

78

Pt

2.2

79

Au

2.4

80

Hg

1.9

81

Tl

1.8

82

Pb

1.8

83

Bi

1.9

84

Po

2.0

85

At

2.2

86

Rn

2.4

87

Fr

0.7

88

Ra

0.9

89

Ac

1.1

104

Rf

-

105

Db

-

106

Sg

-

107

Bh

-

108

Hs

-

109

Mt

-

110

Uun

-

111

Uuu

-

112

Uub

-

113 114

Uuq

-

115 116

Uuh

-

117 118

Uuo

-

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Example 1: What type of bond would form between an atom of nitrogen and an atom of chlorine? a. Nitrogen has an electronegativity value of

3.0.

b. Chlorine has an electronegativity value of

3.0.

c. The difference in the electronegativity values

for nitrogen and chlorine is

ΔEN = - =

d. Therefore the type of bond formed would be

nonpolar covalent. The electrons would be

shared equally. .

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3.0 3.0 0.0

Example 2: What type of bond would form between an atom of hydrogen and an atom of chlorine? a. Hydrogen has an electronegativity value of

2.1.


3.0.


for hydrogen and chlorine is

ΔEN = - =


polar covalent. The electrons would be

shared unequally. .

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3.0 2.1 0.9

Dipole A bond formed between atoms which are not

shared equally is called a dipole.

a) In the bond formed between hydrogen and

chlorine, the chlorine would form the negative

dipole (symbolized by δ-) because it has the

higher electronegativity value.

b) The hydrogen would form the positive dipole

(symbolized by δ+) because it has the lower

electronegativity value.

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Example 3: What type of bond would form between an atom of lithium and an atom of chlorine? a. Lithium has an electronegativity value of

1.0.


3.0.


for lithium and chlorine is

ΔEN = - =


ionic. The electrons would be transferred

between atoms .

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3.0 1.0 2.0

Example 3: What type of bond would form between an atom of lithium and an atom of chlorine?

The lithium atom would lose electrons

and form a positive ion, also known as

a cation.

The chlorine atom would gain electrons

and form a negative ion, also known as

an anion.

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You Try It 1. Complete the following table.

Compound Elements Electronegativity ∆EN Bond Type

KF K

F

O2 O

O

ICl I

Cl

0.8

4.0 3.2 Ionic

3.5

3.5 0.0

Nonpolar

Covalent

2.5

3.0 0.5

Polar

Covalent

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You Try It

2. For each of the bonds in question 1

that were polar covalent, identify the

negative dipole (δ-) and the positive

dipole (δ+).

ICl Iodine is the positive dipole and

chlorine is the negative dipole.

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You Try It

Nonpolar covalent

3. Elements that exist as two atoms

chemically bonded together are called

diatomic elements. The diatomic elements

are hydrogen, bromine, oxygen, nitrogen,

chlorine, iodine, and fluorine. (You need to

memorize the diatomic elements.) What

type of chemical bond exists between the

diatomic elements?

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You Try It

4. Using the three classifications of

bonds discussed, predict the type of

bond that is most likely to be present

in compounds made from elements of

groups 1 (1A) and 17 (7A).

Ionic

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You Try It

5. Using the three classifications of

bonds discussed, predict the type of

bond that is most likely to be present

in compounds made from elements of

groups 16 (6A) and 17 (7A).

Polar Covalent

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You Try It

6. Arrange the following chemical bonds

in order of least covalent to most

covalent: H-H, H-Cl, H-Br, Li-Cl

Li-Cl, H-Cl, H-Br, H-H

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Drawing Lewis Dot Diagrams for Atoms

The electrons that play the most important role in determining whether or not a chemical bond will form are the valence electrons.

In a Lewis dot diagram, dots are placed around the chemical symbol of an element to illustrate the valence electrons. The chemical symbol represents the nucleus of the atom. Back to

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Drawing Lewis Dot Diagrams for Atoms

Examples

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Group 1 Group 2 Group 13 Group 14 Group 15 Group 16 Group 17 Group 18

H He

Li Be B C N O F Ne

Drawing Lewis Structures for Covalent Compounds

Types of Covalent Bonds

Single Covalent Bond – one pair of valence electrons is shared.

Double Covalent Bond - two pairs of valence electrons are shared.

Triple Covalent Bond - three pairs of valence electrons are shared.

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Example 1: H2

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H

H H H

The two hydrogen atoms will

form a single, nonpolar covalent

bond.

Example 2: O2

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O

O O O

The two oxygen atoms will form a

double, nonpolar covalent bond.

Example 3: N2

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N

N N N

The two nitrogen atoms will form

a triple, nonpolar covalent bond.

Example 4: HCl

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H

Cl H Cl

Example 5: NH3

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H

H

N H

You Try It

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HBr CCl4 H2O

C2H5Cl C2H4 C2H2

H2O2 HCN CO2

Structural Formulas Structural formulas can also be

used to show the arrangement of

atoms in molecules.

In a structural formula, dashes are

used to represent shared pairs of

electrons. Back to main menu

Structural Formulas Example: H2

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H H H H –

Structural Formulas Example: H2S

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H S S H –

H H

Structural Formulas Example: N2

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N N N N ≡

Binary Molecular Nomenclature Compounds formed when atoms covalently

bond are called molecular compounds.

Binary molecular compounds are generally

composed of two nonmetallic elements.

When two nonmetallic elements combine,

they often do so in more than one way. For

example carbon can combine with oxygen to

form carbon dioxide, CO2 and carbon

monoxide, CO. Back to main menu

Naming Binary Molecular Compounds

Prefixes are used to show how many atoms

of each element are present in each

molecule of the compound.

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mono- 1

di- 2

tri- 3

tetra- 4

penta- 5

hexa- 6

hepta- 7

octa- 8

nona- 9

deca- 10


The names of molecular

compounds have this

form: (prefix + element

name) (prefix + element

root + ide)

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The prefix mono is usually omitted if there is

just a single atom of the first element.

Example: CO2 is carbon dioxide not

monocarbon dioxide.

If the vowel combinations o-o or a-o appear

next to each other in the name, the first of

the pair is omitted to simplify the name.

Example: N2O is dinitrogen monoxide not dinitrogen monooxide.

You Try It

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Name the following compounds.

a. CBr4

b. Cl2O7

c. N2O5

d. BCl3

e. PCl5

f. NO

a. Carbon tetrabromide b. Dichlorine heptoxide

c. Dinitrogen pentoxide

d. Boron trichloride

e. Phosphorus pentachloride f. nitrogen monoxide

Writing Formulas for Binary Molecular Compounds

To write the formula for a binary molecular

compound you simply write down the

number of atoms of each element indicated

by the name.

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Example: Carbon tetrachloride

CCl4

You Try It

Write formulas for the following binary

molecular compounds.

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a. dinitrogen tetrahydride

b. carbon disulfide

c. iodine heptafluoride

d. sulfur dioxide

N2H4

CS2

IF7

SO2

Writing Formulas for Binary Molecular Compounds

A few molecular compounds have common

names that all scientists use in place of

formal names.

CH4 is methane

H2O is water

NH3 is ammonia

You need to memorize these.

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Exceptions to the Octet Rule

Some molecules are stable even

though the atoms do not all obtain

an octet.

There are three common

exceptions to the octet rule.

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Exception #1

In some molecules the central atom

has less than eight valence electrons.

This is called an incomplete octet.

Incomplete octets are common in

covalent compounds in which the

central atom is beryllium, boron or

aluminum.

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Exception #1

Example: BeH2

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Beryllium has only four electrons

around it.

Exception #2

Molecules almost always have an

even number of electrons, allowing

electrons to be paired, but there

are some exception in which there

are an odd number of electrons.

These exceptions usually involve

nitrogen.

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Exception #2

Example: NO

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You will not be expected to draw

exceptions with odd numbers of

electrons in this course.

Exception #3

In some molecules the central atom has more than eight valence electrons.

This is called an expanded octet.

Some common central elements that have expanded octets are sulfur, chlorine, bromine, iodine, xenon, phosphorus, and arsenic.

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Exception #3

Example: SF6

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Sulfur has twelve electrons around it.

You Try It

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BF3 AsH5 BeI2

PCl5 ClF5 XeF4


Objectives

1. Define coordinate covalent

bonding and give an example.

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A coordinate covalent bond is

formed when one atom contributes

both electrons in a shared pair.

Example: CO

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and Polyatomic Ions

Polyatomic ions form coordinate covalent bonds.

A polyatomic ion is covalently bonded within itself, but is ionically bonded to another atom or polyatomic ion to form a neutral compound.

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and Polyatomic Ions

Example: NH4+

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Example: SO42-

You Try It

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OH- PO43- ClO3

-


Objectives

1. Define resonance and draw

resonance structures for

molecules.

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Resonance occurs when more than

one valid Lewis structure can be

written for a particular molecule.

The different Lewis structures

possible for a molecule are referred

to as resonance structures.

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Let’s look at the Lewis structure for

the ozone, O3, molecule.

Another possible structure for the

ozone molecule is as follows:

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O=O-O:

:O-O=O

. . . .

. . . .

. .

. .

. .

. . . .

. .


Notice that each structure indicates

that the ozone molecule has two

types of O-O bonds, one single

bond and one double bond.

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Based on the Lewis structures we

just drew, you would expect the

bond lengths between the atoms to

be different.

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Scientists, however, have

experimentally determined that the

bond lengths between the oxygen

atoms are identical.

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No one structure correctly describes the ozone molecule. Scientists have determined that the structure for ozone is the average of the two structures.

A double-headed arrow is used to indicate resonance.

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O=O-O: :O-O=O . .

. . . . . .

. .

. . . . . .

. . . .

You Try It

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Resonance structures can often be written

for polyatomic ions. Draw the possible

resonance structures for NO2- .

O=N-O: :O-N=O . .

. .

. . . .

. . . .

. .

. .

. . . .


Objectives

1. Define VSEPR and given a chemical formula of a simple molecule, identify its geometric shape as linear, trigonal planar, angular, tetrahedral, trigonal pyramidal, trigonal bypyramidal, or octahedral.

2. Using the shape of a molecule and electronegativites of its atoms, determine the polarity of the molecule.

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The valence shell electron pair repulsion

(VSEPR) theory can be used to predict the

three dimensional shape of a molecule.

The main idea behind the VSEPR theory is

that electron pairs (bonding and

nonbonding) will orient themselves so that

repulsions between electron pairs are

minimized.

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2 – Atom Linear

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Formula

Lewis

Structure

Drawing

of Model

Bond

Angle

HI H I 180°

Note: There is not a central atom in

a 2-atom linear molecule.

3 – Atom Linear

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Formula Lewis

Structure

Drawing

of Model

Bond

Angle

HCN

HCN is a 3-atom linear molecule. Which atom is the central

atom in the HCN molecule?

How many atoms are bonded to the central atom?

How many pairs of nonbonding electrons on the central

atom of the HCN molecule?

Carbon

None

180° H C N

Two

Bent (also called angular)

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Formula Lewis

Structure

Drawing

of Model

Bond

Angle

H2O 104.5°

Which atom of the water molecule is the central atom?


How many pairs of nonbonding electrons on the central

atom of the H2O molecule?

Oxygen

Two

Two

How can you differentiate between a linear

molecule and a bent molecule in terms of

nonbonding electron pairs on the central

atom?

Linear molecules do not have nonbonding

electrons on the central atom. Bent

molecules have nonbonding electrons on

the central atom.

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Trigonal Planar

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Formula Lewis

Structure

Drawing

of Model

Bond

Angle

H2CO 120°

Which atom is the central atom?


Carbon

Three

How many nonbonded pairs of electrons are there on the

central atom?

Zero

Trigonal Pyramidal

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Formula Lewis

Structure

Drawing

of Model

Bond

Angle

NI3 107°



Nitrogen

Three

How many nonbonded electrons are there on the central

atom?

one

How can you differentiate between a trigonal

planar molecule and a trigonal pyramidal

molecule in terms of nonbonding electron

pairs on the central atom?

Trigonal planar molecules do not have

nonbonding electrons on the central atom.

Trigonal pyramidal molecules have a pair of

nonbonding electrons on the central atom.

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Tetrahedral

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Formula Lewis

Structure

Drawing

of Model

Bond

Angle

CH4 109.5°



Carbon

Four

How many pairs of nonbonded electrons are there on the

central atom?

zero

Trigonal Bipyramidal

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Formula Lewis

Structure

Drawing

of Model

Bond

Angle

PH5

120° 90°



Phosphorus

Five


central atom?

zero

Octahedral

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Formula Lewis

Structure

Drawing

of Model

Bond

Angle

SH6 90°



Sulfur

Six


central atom?

zero

Determining Molecular Polarity

The polarity of each bond, along with the

geometry of the molecule, determines the

polarity of the molecule.

A nonpolar molecule has an even

distribution of molecular charge.

A polar molecule has an uneven distribution

of molecular charge.

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Steps in Determining Molecular Polarity

First determine the geometric shape of the

molecule.

Molecules with nonbonding pairs of

electrons on the central atom are always

polar.

Which two shapes are always polar?

bent and trigonal pyramidal

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If the molecule does not contain nonbonding

pairs of electrons on the central atom, the

polarity is determined by the atoms

surrounding the central atom.

If all of the atoms surrounding the central atom

are the same, the molecule is nonpolar. This

is because the bond dipoles will cancel out.

If all of the atoms surrounding the central atom

are not alike, the molecule is polar. The bond

dipoles will not cancel out. Back to main menu


H-C≡N:

This molecule is polar.

H-Be-H

This molecule is nonpolar.

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You Try It

Determine the polarity of each of the

following molecules.

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a. HI

b. H2O

c. H2CO

d. NI3

e. CH4

polar

polar

polar

polar

nonpolar

Intermolecular Forces of Attraction

Objectives

13. Define van der Waals forces, dipole-dipole forces, hydrogen bonds, and London forces.

14. Given a molecule, identify the dominant type of intermolecular force of attraction.

15. Given chemical formulas for two substances, identify which type of intermolecular forces they exhibit and compare their boiling and freezing points. Back to

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Intramolecular vs. Intermolecular

Intramolecular forces – forces within a molecule that hold atoms together, that is, covalent bonds.

Intermolecular forces – forces between molecules that hold molecules to each other.

These intermolecular forces are collectively referred to as Van der Waals Forces.

They are much weaker than covalent bonds.

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Importance of Intermolecular Forces

The strength of the intermolecular forces can be used to determine whether a covalent compound exists as a solid, liquid, or gas under standard conditions.

Solids have the strongest intermolecular forces of attraction between their particles.

The intermolecular forces of attraction between the molecules of liquids are not as strong as those found between the particles of a solid.

Gases have the weakest intermolecular forces of attraction between their particles.

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Importance of Intermolecular Forces

The strength of the intermolecular forces

can also be used to compare melting and

boiling points.

The more strongly the molecules are

attracted to each other, the higher the boiling

and melting points.

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Types of Intermolecular Forces London Dispersion Forces

London dispersion forces exist in all

covalent molecules, however; they are the

most noticeable between nonpolar

molecules and the nonbonding atoms of

noble gases.

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London dispersion forces arise from the

motion of valence electrons.

From the probability distributions of orbitals,

it is concluded that the electrons are evenly

distributed around the nucleus. However, at

any one instant, the electron cloud may

become distorted as the electrons shift to an

unequal distribution.

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It is during this instant that a molecule develops a temporary dipole.

This temporary dipole introduces a similar response in neighboring molecules, thus producing a short-lived attraction between molecules.

In general the larger the electron cloud, the more likely the molecule is to form temporary dipoles.

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London forces are the weakest type of

intermolecular forces of attraction.

Examples: CO2, H2, Ar

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Types of Intermolecular Forces Dipole-Dipole Forces

Dipole-dipole forces of attraction exist between polar molecules.

Polar molecules contain uneven distributions of charge.

The negative dipole of one molecule is attracted to the positive dipole of another molecule.

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Example of Dipole-Dipole Forces HCl HCl is a polar molecule. The hydrogen end of the molecule forms the positive dipole because it has the lower electronegativity. The chloride end of the molecule forms the negative dipole because it has the higher electronegativity. The chloride end of the molecule is attracted to the hydrogen end of a neighboring molecule.

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H−Cl H−Cl Dipole-dipole forces

Cl−H Cl−H

↓

↑

↓

↑

↓

↑

↓

↑

δ+ δ+

δ+ δ+

δ- δ-

δ- δ-

Types of Intermolecular Forces Dipole-Dipole Forces

Dipole-dipole forces of attraction are

stronger than London dispersion forces.

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Types of Intermolecular Forces Hydrogen Bonding

Hydrogen Bonding is a special type of dipole-dipole

force. Since no electrons are shared or transferred,

hydrogen bonding is not a chemical bond.

Hydrogen bonding exists between where the very

electronegative elements of nitrogen, oxygen and

fluorine are covalently bonded to hydrogen.

Hydrogen bonding occurs between hydrogen and

the unbonded electron pairs of nearby N, O, or F

molecules

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Examples of Hydrogen Bonding

Hydrogen bonding occurs in pure

substances. The hydrogen bonding is

represented by a dotted line.

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Hydrogen bonding can also occur in

mixtures.

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Hydrogen bonding occurs in pure

substances. The hydrogen bonding is

represented by a dotted line.

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Types of Intermolecular Forces

Hydrogen Bonding

Hydrogen bonding is about ten times

stronger than ordinary dipole-dipole forces.

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Identifying the Types of Intermolecular Forces of Attractions

The chart below can help you identify the types of

intermolecular forces of attraction exhibited by a

substance. Reminder: London Dispersion Forces

are exhibited by all covalent molecules.

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You Try It

1. List the intermolecular forces of attraction

in order of increasing strength.

London dispersion forces, dipole-dipole

forces, hydrogen bonding

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You Try It 2. What type of intermolecular forces of attraction would be exhibited by each

of the following substances? Justify your answer. The first one has been

done for you. (Hint: Draw the Lewis Structure for the molecule in order to

help you determine the polarity of the molecule.)

a. NH3

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London dispersion forces, dipole-dipole, hydrogen

bonding. NH3 exhibits London dispersion forces

because all covalent molecules exhibit London

dispersion forces. NH3 exhibits dipole-dipole forces

because it’s a polar molecule. NH3 exhibits

hydrogen bonding because it’s a polar molecule in

which hydrogen is bonded to a nitrogen, oxygen, or

fluorine atom. In this case, hydrogen is bonded to

nitrogen.





b. CO2

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London dispersion forces only

CO2 is a nonpolar molecule. Nonpolar molecules

only exhibit London dispersion forces.





c. HI

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London dispersion forces and dipole-dipole forces

HI exhibits London dispersion forces because all

covalent molecules exhibit London dispersion forces.

HI also exhibits dipole-dipole forces because it’s a

polar molecule. It does not exhibit hydrogen bonding

because since H is not bonded to O, N or F.





d. BeH2

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London dispersion forces only

BeH2 is a nonpolar molecule. Nonpolar molecules

only exhibit London dispersion forces.

Comparing Boiling Points

Two factors that affect boiling point are the

mass of the compound (molar mass) and

the strength of the intermolecular forces of

attraction. The stronger the intermolecular

forces of attraction the higher the boiling

point.

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Examine the table below.

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Boiling Points of Halogens

Name Formula

Physical State at

Room

Temperature

Molar Mass

(g/mol)

Boiling Point

(K, at 1 atm)

fluorine F2 gas 38.0 85.0

chlorine Cl2 gas 70.9 239.1

bromine Br2 liquid 159.8 331.9

iodine I2 solid 253.8 457.4

1. What relationship exists between the mass of

the halogens and the boiling point?

The larger the molar mass, the higher the

boiling point.


Examine the table below.

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Boiling Points of Halogens

Name Formula

Physical State at

Room

Temperature

Molar Mass

(g/mol)

Boiling Point

(K, at 1 atm)

fluorine F2 gas 38.0 85.0

chlorine Cl2 gas 70.9 239.1

bromine Br2 liquid 159.8 331.9

iodine I2 solid 253.8 457.4

2. Arrange the halogens in order of increasing

intermolecular strength of attraction. Justify your answer.

F2, Cl2, Br2, I2

The stronger the intermolecular forces of

attraction, the greater the boiling points.


3. The graph below is a plot of the boiling points of the hydrogen

compounds in the groups headed by fluorine (HF, HCl, HBr, and

HI), oxygen (H2O, H2S, H2Se, H2Te), nitrogen (NH3, PH3, AsH3,

SbH3), and carbon (CH4, SiH4, GeH4, SnH4). Use the graph below

to answer the following questions.

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a. Which group of elements has the

lowest boiling points for each

period? Why do they have the

lowest boiling points for each

period?

The group headed by carbon has

the lowest boiling points for each

period. They are all nonpolar

molecules. Nonpolar molecules

exhibit weaker London dispersion

forces.







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b. Notice in each of the other three groups

that the first compound (H2O, NH3, and

HF) in each group has a significantly

higher boiling point than the other

elements in their groups. What

accounts for this phenomenon?

H2O, NH3, and HF all exhibit hydrogen

bonding. The other substances in the

groups exhibit dipole-dipole forces of

attraction which are not as strong as

hydrogen bonding. Since H2O, NH3, and

HF all exhibit hydrogen bonding they have

higher than expected boiling points.







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c. With the exception of H2O,

NH3, and HF, why do the

boiling points generally

increase within a group?

The boiling points increase

because the molar mass of the

compounds increases.

You Try It

1.Determine whether each of the following

would more likely be formed by polar or

nonpolar molecules.

a. a solid at room temperature

b. a liquid with a high boiling point

c. a gas at room temperature

d. a liquid with a low-boiling point

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polar

polar

nonpolar

nonpolar

You Try It

2. Considering what you have learned about

forces between atoms and molecules,

why do you think all of the elements in

group 18 exist as gases at room

temperature?

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The noble gases exhibit London dispersion

forces. London dispersion forces are the

weakest of the intermolecular forces of

attraction. Substances with weak

intermolecular forces of attraction tend to have

lower boiling points.

You Try It

3. Arrange the following according to

increasing boiling point: H2O, H2S, CO2.

Justify your ranking.

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CO2 < H2S < H2O

CO2 has only London dispersion forces.

H2S has dipole-dipole forces.

H2O has hydrogen bonding.

You Try It

4. Arrange the following according to

increasing boiling point: CH4, CI4, CF4.

Justify your ranking.

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CH4 < CF4 < CI4

All three molecules are nonpolar and thus only

have London dispersion forces between them.

The bigger the molecule, the more electrons

and thus the larger the temporary dipole. The

larger the temporary dipole, the stronger the

intermolecular force and thus the higher the

melting point.

You Try It 5. NH3 is a gas at room temperature and H2O is

a liquid at room temperature. However, they

both exhibit hydrogen bonding. What does

that tell you about the strength of the

hydrogen bonding in H2O as compared to

NH3?

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The hydrogen bonding in H2O is stronger than the hydrogen

bonding which occurs in NH3. H2O has two H atoms that can

potentially form four hydrogen bonds with surround water

molecules. There are exactly the right number of hydrogens and

lone pairs that every one of them can be involved in hydrogen

bonding. In the case of ammonia, the amount of hydrogen

bonding is limited by the fact that each nitrogen has only one lone

pair.

Chemical Bonding (Predicting Bond Types) Lewis (Electron...

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