Chapter 8 Bonding: General Concepts. 8.1 Types of Chemical bonds l A bond is a force that holds atoms together and make them function together l Bond

Chapter 8

Bonding: General Concepts

8.1 Types of Chemical bonds

A bond is a force that holds atoms together and make them function together

Bond energy: the energy required to break a bond.

Why are compounds (atoms aggregate with each other) formed?– Because this situation gives the system

the lowest possible energy.

Ionic Bonding An atom with a low ionization energy

(metal that looses its electron easily) reacts with an atom with high electron affinity (nonmetal).

The electron moves. Opposite charges hold the atoms

together. Closely packed oppositely charged ions

are held by strong electrostatic attraction forces

The highly ordered solid collection of ions is called an ionic compound (crystal)

The net attractive electrostatic forces that hold the cations and anions together are ionic bonds

• Ionic crystals have great thermal stability and consequently acquire high melting points

When atoms lose or gain electrons, they acquire a noble gas configuration, but do not become noble gases

Ionic Bond• When metals and nonmetals

combine, valence electrons usually are transferred from the metal to the nonmetal atoms giving rise to electrostatic attraction force called ionic bond

Energy of interaction between a pair of ions: Coulomb's Law

E= 2.31 x 10-19 J · nm Q is the charge. r is the distance between the centers. If charges are opposite, E is negative Exothermic: the ion pair has less energy the ion pair has less energy

than separated ionsthan separated ions Same charge, positive E, requires energy to

bring them together.

)21

(r

QQ

)21

(r

QQ

What about identical atoms?

The electrons in each atom are attracted to the nucleus of the other.

The electrons repel each other, The nuclei repel each other. They reach a distance with the lowest

possible energy. The distance between is the bond

length.

0

En

erg

y

Internuclear Distance

Atoms are infinitely apart

Zero interacting

energy

How does a bonding force develop

between two identical atoms?

0

En

erg

y


Energy as a function of Internuclear distance

0

En

erg

y


0

En

erg

y


Most stable state

0

En

erg

y


Bond Length

0

En

erg

y


Bond Energy

The molecule is more stable than the two separate atoms

Covalent Bonding Covalent bonding: electrons are

shared equally between two identical atoms.

on the other extreme (ionic bonding) electrons transfer completely to form oppositely charged ions

In between are polar covalent bondspolar covalent bonds. The electrons are not shared evenly. One end is slightly positive, the other

negative. Bond polarity is indicated using small

delta

H - F+ -

Polar covalent bond Polar covalent bond

The density of electron cloud is shifted(some what) towards one of the two bonded atoms

H - F+ -

H - F

+-H - F+

-

H - F

+-

H - F +-

H - F+-

H - F

+-

H - F

+-

H - F+ -

H - F

+-H - F+

-

H - F

+-

H - F +-

H - F+-

H - F

+-

H - F

+-

+-Effe

ct of e

lec

tric field

on

HF

m

ole

cule

s

H - F+ -

H - F+ -

H - F+ - H - F

+ -

H - F+ -

H - F+ -

H - F+ -

H - F+ -

- +

8.2 Electronegativity The ability of an atom in a molecule to attract

shared electrons to itself. To measure the relative electronegativity imagine

an H-X molecule Pauling method: compare the measured

H-X bond energy with the expected H-X bond energy

= (H-X) actual - (H-X)expected

2

energy X-Xenergy H-Henergy X-H Expected

Electronegativity

is known for almost every element Gives us relative electronegativities of all

elements. Tends to increase left to right. decreases as you go down a group. Noble gases aren’t discussed. Difference in electronegativity between

atoms tells about the polarity of the bond

Electronegativity

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

Electronegativity difference & bond type

Electronegativity difference

Bond Type

Zero

Intermediate

Large

Covalent

Polar Covalent

Ionic

Co

valent C

haracter

decreases

Ion

ic Ch

aracter increases

Relationship between electronegativityand bond type

Pauling’s Electronegativities

Electronegativity Differences

8.3 Bond polarity and dipole moments

A molecule with a center of negative charge and a center of positive charge is dipolar (two poles),

or has a dipole moment. Center of charge doesn’t have to be on

an atom. Dipoles will line up in the presence of an

electric field.

Representation of the dipolar character

H - F+ -

• Any diatomic (two atoms) molecule with a polar bond will show a molecular dipole moment

H F

FH

A covalent bond with greater electron density around one of the two atoms

electron richregion

electron poorregion e- riche- poor

+ -

Polar Covalent Bond

Polarity of polyatomic molecules The effect of polar bonds on the polarity of

the entire molecule depends on the molecule shape– carbon dioxide has two polar bonds, and is

linear = nonpolar molecule

Molecules with polar bonds but no resulting dipole moment

Polar molecules The effect of polar bonds on the polarity of

the molecule depends on the molecular shape– water has two polar bonds and a bent

shape; the highly electronegative oxygen pulls the e- away from H = very polar!

Thus, H2O molecule has a dipole moment

How to decide for molecular polarity?

Any diatomic molecule with a polar bond is a polar molecule

For a three or more atoms molecule there are two considerations:

– There must be a polar bond.

– Geometry can’t cancel it out.

Geometry and polarity Three shapes will cancel them out.

Linear

Planar triangles

120º

Tetrahedral

Others don’t cancel, e.g., Bent molecule

8.4 Ions: electron configuration and size

Atoms tend to react to form noble gas configuration.

Nonmetals gain electrons from metals or share electrons with other nonmetals

Metals lose electrons to form cations Nonmetals can share electrons in covalent

bonds. Or they can gain electrons to form anions.

Formation of Cations

Na 1s22s22p63s1 1 valence electron Na1+ 1s22s22p6 This is a noble gas

configuration with 8 electrons in the outer level.

The noble gas configuration (Octet rule)

In forming compounds, atoms tend to achieve a noble gas configuration; 8 e- in the outer level is stable

Each noble gas (except He: 2 e-) has 8 electrons in the outer level

For H it is duet rule (Rule of 2)

Li + F Li+ F -

Ionic bond and the octet rule

1s22s11s22s22p5 1s21s22s22p6[He][Ne]

Li

1s 2s 2p

F

1s 2s 2p

+

Li+

1s 2s 2p

F-

1s 2s 2p+

Noble gas configuration (Octet) rule

and covalent bonding

Atoms are bonded with the electron configuration of a noble gas; that is, the atoms obey the octet rule

EOS

By double-counting the shared electrons in a Lewis structure, each atom appears to have a noble gas configuration

Predicting formulas of Ionic Compounds

When the term “ionic compound” is used; it means solid state (crystalline) of that compound

Ions align themselves to maximize attractions between opposite charges,

and to minimize repulsion between like ions.

Predicting formulas of ionic compounds

Stoichiometry is an important consideration …

MgO

Li2O

1s22s22p63s2 1s22s22p4

Common ions with noble gas configurations in ionic compounds

Predicting formulas

Predict the formula of the compound formed from Al and O

Al : [Ne]3s23p1 (should lose 3e- Ne) O : [He]2s22p4 ) should gain 2e- Ne) The compound to be electrically

neutral it has to be Al2O3

Size of ions

Ion size increases down a group. Cations are smaller than the atoms they

came from. Anions are larger than atoms they came

from. across a row they get smaller, and then

suddenly larger. First half are cations. Second half are anions.

Periodic Trends Across the period, the change is

complicated because of the change from predominantly metals on the left to nonmetals on the right.

Li+

Be2+

B3+

C4+

N3-

O2- F-

Ato

mic an

d Io

nic

Rad

ii

Size of Isoelectronic ions Iso - same Iso electronic ions have the same #

of electrons Al3+ , Mg2+, Na+, Ne, F-, O2- and N3-

All have 10 electrons. All have the configuration 1s22s22p6

Isoelectronic Configurations

Elements that all have the same number of electrons

For isoelectronic species, the greater the nuclear charge, the smaller the species

Effective n

uclear ch

arge

Size of Isoelectronic ions Positive ions have more protons

so they are smaller.

Al+3

Mg+2

Na+1 Ne F-1 O-2 N-3

Ionic solid is formed because the aggregated oppositely charged ions have a lower energy than the original elements

How strongly the ions attract each other in the solid state is expressed by the lattice energy

Lattice energy:

The change in energy that takes place when separated gaseous ions are packed together to form an ionic solid

M+(g) + X-(g) MX

8.5 Energy effects in binary ionic compounds

Lattice energy It is usually defined as the energy

released released when ionic solid is formed from its gaseous ions (it has –ve sign)– M+(g) + X-(g) MX(s)

Lattice energy is a quantitative measure of the stability of ionic compound

The higher the lattice energy the more stable the compound

Na(s) + ½F2(g) NaF(s)

First sublime Na Na(s) Na(g)H = 109 kJ/mol

Ionize Na(g) Na(g) Na+(g) + e-

H = 495 kJ/mol Break F-F Bond ½F2(g) F(g)

H = 77 kJ/mol Add electron to F F(g) + e- F-(g)

H = -328 kJ/mol Formation of NaF from Na+(g) & F-(g) (Lattice energy) Na+(g)+ F-(g) aF(s) H = -1281 kJ/mol

Energy changes associated with the formation of ionic solid

Na(s) + ½F2(g) NaF(s) Lattice energy

Na(s) + ½F2(g) NaF(s)H = -928 kJ/mol

Calculating Lattice Energy

k is a constant that depends on the structure of the crystal electron configuration of ions

r is internuclear distance (r = radius of cation+radius of anion)

Lattice energy is greater with more highly charged ions and distances between ions decrease

)r

QQk(energy Lattice 21

8.6 Partial Ionic Character of covalent bonds

There are probably no totally ionic bonds between individual atoms.

Calculate % ionic character of a bond Compare measured dipole moment of X-

Y bonds to the calculated dipole moment of X+Y- the completely ionic case in the gaseous phase

%100)YX ofmoment dipole calculated

Y-X ofmoment dipole (bond a ofcharacter ionic

_X

measuredPercent

% Io

nic

Ch

arac

ter

Electronegativity difference

25%

50%

75%

The Relationship Between the Ionic Character of a Covalent Bond and the Electronegativity Difference of the Bounded Atoms

What are the ionic compounds?

If bonds can’t be ionic, what are ionic compounds?

What about polyatomic ions?An ionic compound will be defined

as any substance that conducts electricity when melted

8.7 The covalent chemical bond: A model

Bonds are the forces that cause a group of atoms to behave as a unit.

Why do chemical bonds occur?–Due to the tendency of atoms in a system

to achieve its lowest possible energy It takes 1652 kJ to dissociate a mole of CH4

into its separate atoms C & H. Thus, CH4 is stable relative to its separated

atoms

To explain stability, chemical bond term was introduced

Since each hydrogen in CH4 is hooked to the carbon, we get the average energy = 413 kJ/mol = Bond EnergyBond Energy

CH3Cl has 3 C-H, and 1 C – Cl and a stabilization energy of 1578 kJ/mol

Thus, C-Cl bond can be calculated as 339 kJ/mol The bond is a human invention. It is a method of explaining the energy change

associated with forming molecules. Bonds don’t exist in nature, but are useful. We have a model of a bond.

The bond Model

Explains how nature operates.Derived from observations. It simplifies the and categorizes the

information.A model must be sensible, but it

has limitations.

Properties of a Model

Models are human inventions; it does not equal rality.

Models can be wrong, because they are based on speculations and oversimplification.

Become more complicated with age. You must understand the assumptions in the

model, and look for weaknesses. We learn more when the model is wrong than

when it is right.

8.8 Covalent bond energies and chemical reactions

Bond energies (BE) is H when 1 mole of bonds is broken in the gaseous state

H2(g) 2H(g) H = +436 kJ

Cl2(g) 2Cl(g) H = +243 kJ BE is always positive (endothermic) Energy is given off when bonds are formed

H(g) + F(g) HF(g) H = -565 kJ Bond energy increases with bond polarity C-F > N-F > O-F > F-F

Covalent bond energies and chemical reactions

Consider stepwise decomposition of CH4 Each C-H bond has a different energy. CH4 CH3 + H H = 435 kJ/mol

CH3 CH2 + H H = 453 kJ/mol

CH2 CH + H H = 425 kJ/mol CH C + H H = 339 kJ/mol Each bond is sensitive to its environment. Average C-H bond energy = 1652/4=413kJ/mol

Multiple bonds Single bond one pair of electrons is

shared. Double bond two pair of electrons are

shared. triple bond three pair of electrons are

shared. BE is larger for a multiple bond than

for a single bond between the same two atoms

More bonds, shorter bond length.

Av

era

ge

Bo

nd

En

erg

ies

(kj/m

ol)

Bond Type

Bond Length

)pm(

C-C 154

CC 133

CC 120

C-N 143

CN 138

CN 116

Lengths of Covalent Bonds

Bond Lengths

Triple bond < Double Bond < Single Bond

Bond Energies and Enthalpy changes in reactions

H = D required– D released when BE bonds broken BE bonds formed

= BE (reactants) – SBE (products)

9.10

Enthalpy change for a reaction

Use bond energies to calculate the enthalpy change for:H2 (g) + F2 (g) 2HF (g)

H = BE(reactants) – BE(products)

Type of bonds broken

Number of bonds broken

Bond energy (kJ/mol)

Energy change (kJ)

H H 1 436.4 436.4

F F 1 156.9 156.9

Type of bonds formed

Number of bonds formed

Bond energy (kJ/mol)

Energy change (kJ)

H F 2 568.2 1136.4

H = 436.4 + 156.9 – 2 x 568.2 = -543.1 kJ

Find the energy for this

2 CH2 = CHCH3

+

2NH3 O2+

2 CH2 = CHC N

+

6 H2O

C-H 413 kJ/molC=C 614kJ/molN-H 391 kJ/mol

O-H 467 kJ/molO=O 495 kJ/molCN 891 kJ/mol

C-C 347 kJ/mol

8.9 The localized electron bonding model

Simple model, easily applied to describe covalent bonds

A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.

Electron pairs are assumed to be localized on a particular atom (Lone pairs) or in the space between two atoms (Bonding pairs)

The Localized Electron Model

has three parts: 1. Description of valence electrons

arrangements in the molecule using Lewis structures

2. Prediction of geometry of the molecule using VSEPR (valence shell electron- pair repulsion model

3. Description of the types of orbitals used by the atoms to share electrons or hold lone pairs.

8.10 Lewis Structure

Shows how the valence electrons are arranged around atoms in a molecule

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.

Nitrogen, N, is in Group 5A and therefore has 5 valence electrons.

N:.

..

:

N .. ..N :.

. :N ...

Place one dot per valence electron on each of the four sides of the element symbol.

Pair the dots (electrons) until all of the valence electrons are used.

Lewis Dot Symbols

1. Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center.

2. Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge.

3. Use one pair of electrons to form a bond (a single line) between each pair of atoms.

4. Arrange the remaining electrons to satisfy an octet for all atoms (duet for H), starting from outer atoms.

5. If a central atom does not have an octet, move in lone pairs to form double or triple bonds on the central atom as needed.

Rules for Writing Lewis Structures

Examples Fluorine has seven valence electrons

9F: 1s22s22p5

A second atom also has seven By sharing electrons… …both end with full orbitals

F F8 Valence electrons

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

Write the Lewis structure of nitrogen trifluoride (NF3).

Step 1 – N is less electronegative than F, put N in center

F N F

F

Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)

5 + (3 x 7) = 26 valence electrons

Step 3 – Draw single bonds between N and F atoms.

Step 4 – Arrange remaining 20 electrons to complete octets

Write the Lewis structure of the carbonate ion (CO32-).

Step 1 – C is less electronegative than O, put C in center

O C O

O

Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4) -2 charge – 2e-

4 + (3 x 6) + 2 = 24 valence electrons

Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms.

Step 4 - Arrange remaining 18 electrons to complete octets

Step 5 – The central C has only 6 electrons. Form a double bond.

2

Writing Lewis Structures

If you run out of electrons before the central atom

has an octet…

…form multiple bonds until it does.

Maximum number of bonds (or atoms) surrounding the central atom

Central atomMax # bondsH1H-O2 -O-

N3 -N-

C4 -C-

X-(F, Cl, Br, I)1X-

Exceptions to the Octet Rule

The Incomplete Octet

H HBeBe – 2e-

2H – 2x1e-

4e-

BeH2

BF3

B – 3e-

3F – 3x7e-

24e-

F B F

F

3 single bonds (3x2) = 69 lone pairs (9x2) = 18

Total = 24

The expanded octet

Some compounds have expanded valence shells, which means that the central atom has more than eight electrons around it

#Ve= 5 + 7(5)= 40 #Ve = 6 + 6(7)= 48

Expanded octet

EOS

An expanded valence shell may also need to accommodate lone-pair electrons as well as bonding pairs

Molecules with expanded octet Terminal atoms are most often halogens. In few

cases one or more O-atoms are at the end of the molecule

The central atom is a nonmetal in the 3rd, 4th or 5th period of the Periodic Table– 3rd P S Cl– 4th As Se Br Kr– 5th Sb Te I Xe

All atoms have d-orbitals available for bondig (3d, 4d, 5d) where the extra pairs of electrons are located

Elements of Period 2, NEVERPeriod 2, NEVER form compounds with expanded octet.

Odd-Electron Molecules

N – 5e-

O – 6e-

11e-

NO N O

N O

Species with unpaired electrons show weak attraction to magnetic field

A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure.

O O O OOO

O C O

O

- -O C O

O

-

-

OCO

O

-

- 9.8

What are the resonance structures of the carbonate (CO3

2-) ion?

8.12 Resonance

Resonance•In truth, the electrons that form the second C—O bond

in the double bonds below do not always sit between that C and that O, but rather can move among the two

oxygens and the carbon.•They are not localized, but rather are delocalized.

Molecules that don’t follow the octet rule

EOS

Molecules with an odd number of valence electrons have at least one of them unpaired and are called free radicalsfree radicals

Formal Charge

For molecules and polyatomic ions that exceed the octet there are several exceed the octet there are several different structuresdifferent structures.

Use charges on atoms to help decide which one is the real molecule.

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

Formal chargeFormal charge is the difference between the number of valence electrons on the free atom and that assigned to the atom in the molecule.

CCff = N = Nvv – (N – (Nuu + ½ N + ½ Nbb))– Cf = formal charge

– Nv = #valence e- in the un-bonded atom

– Nu = # unshared e- owned by the atom

– Nb= # bonding e- shared by the atom

In molecules Cf is close to zero In ions should be equal to the charge on the ion

Calculation of Formal Charge

Using the assumption of formal charges to evaluate Lewis structure

Atoms in molecules try to achieve as low a formal charge (as close to zero) as possible

Negative formal charges are expected to be found on the most electronegative elements.

Which is the most likely Lewis structure for CH2O?

9.7

H C O H

-1 +1 HC O

H

0 0

8.13 Molecular Structure: VSEPR

Lewis structures tell us how the atoms are connected to each other.

They don’t tell us anything about shape. The shape of a molecule can greatly

affect its properties. VValencealence SShell hell EElectronlectron PPair air RRepulsionepulsion

Theory allows us to predict geometry

VSEPRMolecules take a shape that puts

electron pairs as far awayfar away from each other as possible.

The electron-pairs surrounding an atom (valence electrons) repel one another and are oriented as far apart as possible

Structure around a given atom is determined pricipally by minimizing electron –pair repulsion

To determine electron pairs Lewis structure should be drawn

Find bonding and nonbonding lone pairs

Lone pair take more space. Multiple bonds count as one pair.

VSEPR The number of pairs determines

–bond angles–underlying structure

The number of atoms determinesThe number of atoms determines –actual shape

Valence shell electron pair repulsion (VSEPR) model

Predict the geometry of the molecule from the electrostatic repulsions between the electron (bonding and nonbonding) pairs.

AB2 2 0

Class

# of atomsbonded to

central atom

# lonepairs on

central atomArrangement of electron pairs

MolecularGeometry

10.1

linear linear

B B

Cl ClBe

2 atoms bonded to central atom0 lone pairs on central atom 10.1

AB2 2 0 linear linear

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR

AB3 3 0trigonal planar

trigonal planar

F

F F

B


Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR


trigonal planar

10.1

AB4 4 0 tetrahedral tetrahedral

H

H

H

H

C


Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR


trigonal planar


AB5 5 0trigonal

bipyramidaltrigonal

bipyramidal

P

Cl

Cl

Cl

Cl

Cl


Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR


trigonal planar


AB5 5 0trigonal

bipyramidaltrigonal

bipyramidal

AB6 6 0 octahedraloctahedral

S

F

F

F

F

F

F

Effect of lone pairs on Geometry

• Unshared pair of electrons (under the influence of one nucleus) spreads out over a volume larger than a bonding pair (under the influence of two nuclei).

• The electron pair geometry is approximately same as that observed when only single bonds are involved

• The bond angles are either equal to the ideal values or little less

• The molecular geometry is quite different when line pairs are involved.

• Molecular geometry refers only to the positions of the bonded atoms

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR


trigonal planar

AB2E 2 1trigonal planar

bent

< 120o

SO2

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR

AB3E 3 1


tetrahedraltrigonal

pyramidal

< 109.5o

107o

NH3

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR


AB3E 3 1 tetrahedraltrigonal

pyramidal

AB2E2 2 2 tetrahedral bent

H

O

H

< 109.5o

104.5o

H2O

ABE31 3 LinearH-B

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR

AB5 5 0trigonal

bipyramidaltrigonal

bipyramidal

AB4E 4 1trigonal

bipyramidaldistorted

tetrahedron

SF4

Seesaw90o, 120o, 180o

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR

AB5 5 0trigonal

bipyramidaltrigonal

bipyramidal

AB4E 4 1trigonal


tetrahedron

AB3E2 3 2trigonal

bipyramidalT-shaped

ClF

F

FClF3

90o, 180o

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR

AB5 5 0trigonal

bipyramidaltrigonal

bipyramidal

AB4E 4 1trigonal


tetrahedron

AB3E2 3 2trigonal

bipyramidalT-shaped

AB2E3 2 3trigonal

bipyramidallinear

I

I

II3 180o

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR


AB5E 5 1 octahedral square pyramidal

Br

F F

FF

F

BrF5

90o, 180o

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR


AB5E 5 1 octahedral square pyramidal

AB4E2 4 2 octahedral square planar

Xe

F F

FFXeF4

90o, 180o

10.1

The VSEPR model and multiple bond

• For the geometry purposes:

A multiple bond behaves exactly as if it were a single electron-pair

Molecules containing no single central atom

• The central atoms of the molecule should be labeled first.

• Geometry can be predicted by focusing on each central atom by counting the electron pairs around each central atom.

END

Predicting Molecular Geometry1. Draw Lewis structure for molecule.

2. Count number of lone pairs on the central atom and number of atoms bonded to the central atom.

3. Use VSEPR to predict the geometry of the molecule.

What are the molecular geometries of SO2 and SF4?

SO O

AB2E

bent

S

F

F

F F

AB4E

distortedtetrahedron

10.1

Dipole Moments and Polar Molecules

10.2

H F

electron richregion

electron poorregion

= Q x rQ is the charge

r is the distance between charges

1 D = 3.36 x 10-30 C m

10.2

10.2

10.2

Which of the following molecules have a dipole moment?H2O, CO2, SO2, and CH4

O HH

dipole momentpolar molecule

SO

O

CO O

no dipole momentnonpolar molecule

dipole momentpolar molecule

C

H

H

HH

no dipole momentnonpolar molecule

Does CH2Cl2 have a dipole moment?

10.2

10.2

How well does it work? Does an outstanding job for such a

simple model. Predictions are almost always

accurate. Like all simple models, it has

exceptions.

Documents

Chapter 8 Bonding: General Concepts. 8.1 Types of Chemical bonds l A bond is a force that holds atoms together and make them function together l Bond