Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Bonds Forces that hold groups of...

Copyright©2000 by Houghton Mifflin Company. All rights reserved.

1

Bonds

Forces that hold groups of atoms together and make them function as a unit.


2

Bond Energy

It is the energy required to break a bond.

It gives us information about the strength of a bonding interaction.


3

Bond Length

The distance where the system energy is a minimum.


4

Ionic Bonds

Formed from electrostatic attractions of closely packed, oppositely charged ions.

Formed when an atom that easily loses electrons reacts with one that has a high electron affinity.


5

Ionic Bonds

Q1 and Q2 = numerical ion charges

r = distance between ion centers (in nm)

E = 2.31 10 J nm (19 QQ r1 2 / )


6

Electronegativity

The ability of an atom in a molecule to attract shared electrons to itself.

= (H X)actual (H X)expected


7

Polarity

A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment.

+

FH


8

Achieving Noble Gas Electron Configurations (NGEC)

Two nonmetals react: They share electrons to achieve NGEC.

A nonmetal and a representative group metal react (ionic compound): The valence orbitals of the metal are emptied to achieve NGEC. The valence electron configuration of the nonmetal achieves NGEC.


9

Isoelectronic Ions

Ions containing the the same number of electrons

(O2, F, Na+, Mg2+, Al3+)

O2> F > Na+ > Mg2+ > Al3+

largest smallest


10

Lattice Energy

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

M+(g) + X(g) MX(s)

Lattice energy is negative (exothermic) from the point of view of the system.


11

Formation of an Ionic Solid

1. Sublimation of the solid metal

M(s) M(g) [endothermic]

2. Ionization of the metal atoms

M(g) M+(g) + e [endothermic]

3. Dissociation of the nonmetal 1/2X2(g) X(g) [endothermic]


12

Formation of an Ionic Solid(continued)

4. Formation of X ions in the gas phase:

X(g) + e X(g) [exothermic]

5. Formation of the solid MX

M+(g) + X(g) MX(s) [quite exothermic]


13

Lattice Energy = k( / )QQ r1 2

Q1, Q2 = charges on the ions

r = shortest distance between centers of the cations and anions


14

Models

Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.


15

Fundamental Properties of Models

A model does not equal reality.

Models are oversimplifications, and are therefore often wrong.

Models become more complicated as they age.

We must understand the underlying assumptions in a model so that we don’t misuse it.


16

Bond EnergiesBond breaking requires energy (endothermic).

Bond formation releases energy (exothermic).

H = D(bonds broken) D(bonds formed)

energy required energy released


17

Localized Electron Model

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


18

Localized Electron Model

1. Description of valence electron arrangement (Lewis structure).

2. Prediction of geometry (VSEPR model).

3. Description of atomic orbital types used to share electrons or hold long pairs.


19

Lewis Structure

Shows how valence electrons are arranged among atoms in a molecule.

Reflects central idea that stability of a compound relates to noble gas electron configuration.


20

Comments About the Octet Rule

2nd row elements C, N, O, F observe the octet rule.

2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive.

3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals.

When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

21

Writing Lewis Formulas:The Octet Rule

N - A = S rule Simple mathematical relationship to help us write Lewis dot

formulas. N = number of electrons needed to achieve a noble gas

configuration. N usually has a value of 8 for representative elements. N has a value of 2 for H atoms.

A = number of electrons available in valence shells of the atoms. A is equal to the periodic group number for each element. A is equal to 8 for the noble gases.

S = number of electrons shared in bonds. A-S = number of electrons in unshared, lone, pairs.

22


For ions we must adjust the number of electrons available, A. Add one e- to A for each negative charge. Subtract one e- from A for each positive charge.

The central atom in a molecule or polyatomic ion is determined by: The atom that requires the largest number of electrons

to complete its octet goes in the center. For two atoms in the same periodic group, the less

electronegative element goes in the center.

23

Writing Lewis Formulas:The Octet Rule Example 7-2: Write Lewis dot and dash

formulas for hydrogen cyanide, HCN. N = 2 (H) + 8 (C) + 8 (N) = 18 A = 1 (H) + 4 (C) + 5 (N) = 10 S = 8 A-S = 2 This molecule has 8 electrons in shared

pairs and 2 electrons in lone pairs.

H C N·· ·· ···· H C N ··or··

24


Example 7-3: Write Lewis dot and dash formulas for the sulfite ion, SO3

2-. N = 8 (S) + 3 x 8 (O) = 32 A = 6 (S) + 3 x 6 (O) + 2 (- charge) = 26 S =

6 A-S = 20 Thus this polyatomic ion has 6 electrons in

shared pairs and 20 electrons in lone pairs. Which atom is the central atom in this ion?

You do it!

25


What kind of covalent bonds, single, double, or triple, must this ion have so that the six shared electrons are used to attach the three O atoms to the S atom?

O S O

O··

····

····

··

··

··

····

··

····

2-O S

O

O·· ·· ··

······ ··

······

2-or

LEWIS STRUCTURES

1 : draw skeleton of species2 : count e- in species3 : subtract 2 e- for each

bond in skeleton4 : distribute remaining e-


27

Resonance

Occurs when more than one valid Lewis structure can be written for a particular molecule.

These are resonance structures. The actual structure is an average of the resonance structures.

28

Resonance

Example 7-4: Write Lewis dot and dash formulas for sulfur trioxide, SO3.

You do it! N = 8 (S) + 3 x 8 (O) = 32 A = 6 (S) + 3 x 6 (O) = 24 S = 8 A-S = 16

orO S O

O··

····

····

··

····

····

·· ·· O S

O

O·· ······ ··

······

29

Resonance

There are three possible structures for SO3. The double bond can be placed in one of three places.

O S

O

O·· ······ ··

······

OS

O

O·· ···· ·· ··

··

······

O S

O

O·· ····

·· ··

····

oWhen two or more Lewis formulas are necessary to show the bonding in a molecule, we must use equivalent resonance structures to show the molecule’s structure.o Double-headed arrows are used to indicate resonance formulas.

30

Resonance

Resonance is a flawed method of representing molecules.There are no single or double bonds in

SO3.

• In fact, all of the bonds in SO3 are equivalent.

The best Lewis formula of SO3 that can be drawn is: SO O

O

RESONANCE

a condition occurring when more than one valid Lewis structure can be written for a particular species

RESONANCEThe actual electronic

structure is not represented by any one of the Lewis structures; it is an average of them

33

Writing Lewis Formulas:Limitations of the Octet Rule

There are some molecules that violate the octet rule. For these molecules the N - A = S rule does not apply:

1. The covalent compounds of Be.

2. The covalent compounds of the IIIA Group.

3. Species which contain an odd number of electrons.

4. Species in which the central element must have a share of more than 8 valence electrons to accommodate all of the substituents.

5. Compounds of the d- and f-transition metals.

34


In those cases where the octet rule does not apply, the substituents attached to the central atom nearly always attain noble gas configurations.

The central atom does not have a noble gas configuration but may have fewer than 8 (exceptions 1, 2, & 3) or more than 8 (exceptions 4 & 5).

35

Writing Lewis Formulas:Limitations of the Octet Rule Example 7-5: Write dot and dash

formulas for BBr3.This is an example of exception #2.

You do it!

B··. Br··

··

··.

BBr Br

Br

····

····

····

····

····

····

Br B

Br

Br··

····

·· ····

····

··

or

36


Example 7-6: Write dot and dash formulas for AsF

5.

You do it!

As··

..

. F····

··.

··

As

F

F F

F F

····

··

·· ····

····

····

··

··

·· ··or

····

····

··

·· ····

····

····

··

·· AsF

F F

FF

······ ··

··

··


37

VSEPR Model

The structure around a given atom is determined principally by minimizing electron pair repulsions.


38

Predicting a VSEPR Structure

1. Draw Lewis structure.

2. Put pairs as far apart as possible.

3. Determine positions of atoms from the way electron pairs are shared.

4. Determine the name of molecular structure from positions of the

atoms.

Distinguish between ELECTRONIC

Geometry&

MOLECULARGeometry

CH4

Electronic geometry: tetrahedral

Molecular geometry: tetrahedral

Bond angle = 109.50

H3O+


Molecular geometry: trigonal pyramidal

Bond angle ~ 1070

H2O


Molecular geometry: bent

Bond angle ~ 104.50

NH2-


Molecular geometry: bent

Bond angle ~ 104.50

“Octet Rule” holds for connecting atoms, but may not for the central

atom.

BF3

Electronic geometry: trigonal planar

Molecular geometry: trigonal planar

Bond angle =1200

PF5

Electronic geometry: trigonal bipyramidal

Molecular geometry: trigonal bipyramidal

Bond angle = 1200 & 900

SF4Electronic geometry: trigonal bipyramidal

Molecular geometry: see-saw

Bond angle = 1200 & 900

ICl3

Electronic geometry: trigonal bipyramidal

Molecular geometry: T-shape

Bond angle <= 900

I3-

Electronic geometry: trigonal bipyramid

Molecular geometry: linear

Bond angle = 1800

PCl6-

Electronic geometry: octahedral

Molecular geometry: octahedral

Bond angle = 900

BrF5


Molecular geometry: square pyramidal

Bond angle ~ 900

ICl4-


Molecular geometry: square planar

Bond angle = 900

What happens when there is not enough

electrons to “satisfy” the central atom?Double or Triple

Bonds are needed.

EXAMPLES

EtheneAcetic AcidOxygenNitrogen


55

Formal Charge

The difference between the number of valence electrons (VE) on the free atom and the number assigned to the atom in the molecule.

We need:

1. # VE on free neutral atom

2. # VE “belonging” to the atom in the

molecule

FORMAL CHARGE

the charge assigned to an atom in a

molecule or polyatomic ion

Formal charge =

#VE on free atom -

#VE assigned to atom

in Lewis Structure

Assignment of e-

1 1. Lone pairs belong entirely to atom in question

2 2. Shared e- are divided equally between the two sharing atoms


59

Formal Charge

Not as good Better

CO O(-1) (0) (+1)

CO O(0) (0) (0)

EXAMPLES

Nitrate ionOzone

The sum of the formal charges of all atoms in a species must equal the overall charge on

the species.

If nonequivalent Lewis structures exist, the

one(s) that best describe(s) the bonding

in the species has...

FAVORED LEWIS STRUCTURES

1 1. formal charges closest to zero

2 2. negative formal charge is on the most electronegative atom

EXAMPLES

Carbon dioxideThiocyanate ionSulfate ion

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Bonds Forces that hold groups of...

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