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CM1502

Chapter 3

1CM1502 Sem2-2013-14

Models ofchemical bonding

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Atomic properties and chemical bonds

Why do atoms bond at all?

Properties of an atom- Electronic config./ Zeff

Properties of a substance-Type/Strength of chemical bonds

less stable

more stable

P.E. = -K.E.

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Types of Bonding

Along the line IEs are about 8 eV

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Types of Bonding1. Metal with nonmetal:

ionic bonding

Metal loses electrons to form a positive ion (cation).

Non metal gains electrons to form a negative ion (anion).The electrostatic attraction between the ions draw them

into a three dimensional array to form an ionic solid.

Chemical formula is the empirical formula.

Movie (IVLE workbin-videos-Formation of ioniccompound)

2. Nonmetal with nonmetal:

covalent bonding

The atoms are drawn together as the nucleus of each

atom attracts the electrons of other.The electrons are shared.

Shared pair of electrons is localized between the atoms.

Chemical formula is the molecular formula

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Types of bonds

3. Metal with metal:

electron pooling and metallic bonding

Outer electrons of metals are losely

held due toshielding.These electrons move freely through

the entire piece of metal.

Hence electrons in metallic bonding

are delocalized.

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Lewis electron dot symbolSteps to write the Lewis symbol for main group elements:

1. Note the group number and is the same as their valence electrons.

2. Place one dot at a time on each of the four sides of the symbol.

3. Keep adding dots, pairing them until all are used up.

Octet rule for representative elements: When atoms bond they lose, gain or share

electrons to attain a filled outer level of eight electrons

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7

The Ionic Bonding Model

Lithium Fluoride: The three ways to depict the electron transfer

Electron configurations

Li 1s22s1

Orbital diagrams

Lewis electron-dot symbols

+ F 1s22s22p5 Li+ 1s2 + F 1s22s22p6

Li1s 2s 2p

F

1s 2s 2p

+

Li+

1s 2s 2p

F-

1s 2s 2p

+

.

+ F: ::

Li . Li+ + F::

:

:

Problem: Depict the formation of Potassium Oxide and Barium chlorideCM1502 Sem2-2013-14

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Ionic Bonding Begin with a neutral Cl

atom and a neutral Naatom.

Eionis the energy price

to pay to convert both intoits respective noble gasconfiguration.

The cation and anionattract each other, so atsome distance between

Na+and Cl-the potentialenergy of attraction winsover the positive valueEion.

The cation and anion repeleach other when close

because the e clouds ofboth ions can notinterpenetrate due to thePauli exclusion principle.

It is the strong Coulombattraction that binds allsalts

together.

Eion= IE(Na) EA(Cl)

Na Na++ e Cl- Cl + e

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Lattice energy

Li(s) Li(g) H =161 kJ

F2 F(g)

H = 79.5 kJLi(g) Li+

(g)+e- IE1= 520kJ

F(g)+e- F-(g) EA = -328 kJ

The electron transfer

process actually absorbsenergy !

Generally formation of ionic

solids releases energy.

Li+

(g)+ F-

(g) LiF(s) H0

= -1050kJ

The energy released when an

ionic solid is formed from its

ions is called Lattice energy.

Hence the overall reaction Li(s)+ F2 LIF(s)H = -617.5kJ; exothermic

Ionic solids exists only

because the lattice energy

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Periodic trends in Lattice energy

The lattice energy of L iF and MgO are in the ratio 1:4. Why?

Lattice energy = kQ1Q2

r

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k is the proportionality constant

Q1Q2are the charges on the ions

R is the shortest distance between thecenters of cation and anion

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Explaining the properties A typical ionic compound is hard and brittle.

This is because there is strong attractive forces that hold the ions

together.

They have high meltingand boiling points

They do not conduct electricity in solidstate but do so when melted or

dissolved.

This is because the ions can move in the molten or dissolved stateCompound mp (0C) bp (0C)

CsBr

661

1300

NaI

MgCl2

KBr

CaCl2

NaCl

LiF

KF

MgO

636

714

734

782

801

845

858

2852

1304

1412

1435

>1600

1413

1676

1505

3600

Solid ionic

compound

Molten ionic

compound

Ionic compound

dissolved in

water

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Why do ionic solids crack

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The covalent bonding model

The H2molecule

As the two hydrogen atoms are brought together,

1. The electrons in the two atoms repeleach other

because they have the same charge (E > 0);2. The protons in adjacent atoms repeleach other (E > 0);

3. The electron in one atom is attracted to the oppositely

charged proton in the other atom, and vice versa (E < 0);

At the observed bond distance the repulsive and attractiveinteractions are balanced

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Covalent bond formation in H2

Movie: IVLE-work bin-videos-covalent bond_2

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Properties of covalent bondBond order: is the number of electron pairs being shared by a

given pair of atoms.

Bond energy: Also called Bond enthalpy or bond

strength. It is defined as the standard enthalpy

change for breaking the bond in 1 mol of gaseousmolecules.

A-B (g) A(g)+ B(g) H = BEA-B(always >0)

Stronger bonds have higher bond energy and

weaker ones have lower bond energy

Bond length: The distance between the nuclei of the

bonded atoms

H H: H F: :::

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Relationship between bond order, bond length and bond energy.

For a given pair of atoms, a higher bond order results in shorter bond length

and higher bond energy

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Most covalent substances melt and boil at low temperatures

Pentane

The two forces in covalent molecules are

1. Strong bonding forces that hold the atoms together with the molecule and

2. Weak intermolecular forces that act between the molecules.

When the covalent liquid boils it is the weak forces between the molecules are

overcome and not the strong bonds within the molecule. Hence it needs less energy.

Explaining the Physical properties

Network covalent solids have exceptional behavior

Melts at 1550C Melts at 3550C

Most of the covalent substances are poor electrical conductors.

Reason: 1. electrons are localized

2. No ions are present

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Bond energy and enthalpy.

The heat released or absorbed during a chemical change is due

to differences between reactant and product bond energies.Horeaction = H

obonds broken+ H

obonds formed

Horeaction = reactant bonds broken- productbonds formed

H2(g) + F2(g) 2HF(g) Hreaction = -546kJ

Horeaction = [1x H-H + 1xF-F] [2xH-F]

=[432 + 159] [2x565]

= -539 kJ

Test your self:Calculate theHorxn for chlorination of methane to form chloroform.

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Bond energies and food/fuel

A fuel generally consist of C-H,C-C C-O and O-H bonds.

Of which C-H and C-C bonds are weakerand C-O and O-H bonds are stronger.

Generally a fuel reacts with O2, all the bonds break and form C=O and O-H bonds.

If a fuel consist of many C-C and C-H bonds and fewer C-O , O-H bonds, said to

release Higher energy and these are known as good fuels.

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20

Between the two extremes

EN is the relative ability of a bonded atom to attract shared electrons.

432kJ/mol

H-H + F-F 2 H-F

159kJ/mol Expected BE of H-F is 296kJ/mol

Actual BE of H-F is 565kJ/mol

The reason for increased BE is Electrostatic attraction

Arbitrary cutoff divides ionicfrom covalent bonds

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Polar covalent bonds

Figure 9.23

EN

3.0

2.0

0.0

Electrons are not transferred

completely.

Electrons are not shared

equally.

One atom has a stronger

attraction for the sharedelectron than the other atom.

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Bond polarity and dipole moment

A molecule that has a positive center of charge of

magnitude Q and a negative center of charge ofmagnitude Q separated by a distance R has a

dipole moment() of QR

Dipole moment = = QR

SI Unit = C m (coulomb meter)

Often used unit = debye

1debye(D) = 3.336 x 10-30Cm)

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Molecules with polar bonds and a dipole moment

Molecules with polar bonds and no resulting dipole moment
http://www.cengage.com/chemistry/book_content/9781111580650_zumdahl/images/ch13/13p604_f05c.htmlhttp://www.cengage.com/chemistry/book_content/9781111580650_zumdahl/images/ch13/13p604_f05b.html

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Metallic bonding

All the metal atoms in the sample contribute their valenceelectrons to form a delocalized electron sea.

The piece is held together by the mutual attractionof themetal cations and the mobile electrons.

The metal ion array is regular but not rigid.

Not rigid hencecannot come under ionic bonding

No localized sharing - hence cannot be covalent bond.

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Properties of metalsMelting points are only moderately high.

Reason: the cations can move without

breaking the attraction to thesurrounding

Boiling points are very high

Reason: Higher energy is needed to break

the cation from all the valence electrons.

Periodic trends:

M.Pt decrease down the group.

Reason: Larger metal ions have a weaker

attraction to the electron sea.

M.Pt increase across the period.

Reason: The charge on the cation increases

from left to right, hence stronger attraction

towards the electron sea.

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Why do metals dent and bend?

When hammered, the metalsslide past each other through

the electron sea and end up in

new positions.

Properties of metals

Metals are good electrical and thermal conductorsin both

solid and liquid states because of their mobile electrons.

Presence of foreign atoms disrupt the array and reduce the

conductivity.

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Summary

Atomic properties and type of bond

Features of ionic bonding

lattice energy

properties-ionic compounds are brittle, high melting, conduct

electricity only in molten/dissolved state.

Features of covalent bonding

non metals

bond order bond energy and bond length

Polar covalent compounds

Features of metallic bonding

electron sea metals bend, have high melting and boiling points and conduct

electricity.

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Role of molecular shape

Interactions of reactants during a chemical reaction isbased on the molecular shapes.

Predict the physical and chemical behavior of syntheticmaterials.

Molecular shape is a crucial property of living systems

Eg: Hormonal regulation and function of genes.

Sugar Enzyme 28CM1502 Sem2-2013-14

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What is Lewis Structure?

Two-dimensional structural formula consists

of electron-dot symbols.

It shows which are the atoms bonded to

each other but it does NOT indicate

the three-dimensional shape.

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Molecular formula to Lewis structure.

Molecularformula

Atom

placement

Sum of

valence e

Remaining

valence e

Lewis

structure

Place atom with lowest

EN in center

Draw single bonds. Subtract2e for each bond.

Give each atom 8e

(2efor H)

Step 1

Step 2

Step 3

Step 4

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Molecular

formula

Atom

placement

Sum of

valence e-

Remaining

valence e-

Lewis

structure

NF3

N

FF

F

N 5e-

F 7e- X 3 = 21e-

Total 26e-

:

: :

::: :

:

:

:

31

Electrons involved in bonding are called bond

pairs . These are shared between atoms.

Unshared electrons are called lone pairs .They

belong to only one atom.

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L i St t 1 0

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Lewis Structures v1.0

An atoms valency should be satisfied.

N2must be a triple bond in order to complete theoctet.

O2a double bond in order to complete the octet. F2a single bond in order to complete the octet.

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More Lewis Structures

Note the first period nonmetal (H) completedshell is 2 electrons [He].

CO2can be readily explained. CO, however, can satisfy the octets, but the

valencyof C?

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Valency Rule CAN be Broken

Octets all completed,

valencies all satisfied

Octets all completed,

but C and O valency

broken.

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Alert! Issue Formal ChargePatch.


For nonmetals, when an atoms valency is broken, and theoctet rule is satisfied, the atom must possess either feweror greater electrons than it has valence electrons.

The difference is accounted for by placing a formal chargeon the offending atoms.

So in

This is because C owns 5 electrons in the above

structure but should own 4, and O owns 5 electrons, butshould own 6.

movie (IVLE /workbin/videos/Formal charge calculation)

- +

35

Formal charge of an atom =

No. of valence e- no. of lone pair e- no. of bonding e-

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Octets Not Always Completed

If a molecule has an odd number of electrons,there is no way the octet rule can hold, or thebond order (BO) can be an integer.

NO (nitric oxide) is an example.

Here the BO is 2 and is consistent

with the bond dissociation energy (De)of 6.52 eV

H2+also exists, and must

have a BO of .We will also see later that

He2+also exists and has a

BO also of !

Octet satisfiedOctet notsatisfied

Octets satisfied

BO is 2 But need weird

formal charges

-+

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Lewis Structures with Multiple Bonds.

Step 5: Itfollows the other steps in Lewis structure

construction. If a central atom does not have 8e-, an octet,then a lone pair of e-can be moved to form a multiple bond.

CCH

H H

H

37

CCH

H H

H

:

H C C H

37CM1502 Sem2-2013-14

E di th O t t

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Expanding the Octet Unless required, or stated otherwise, we always stick to the

octet rule.

We are forced in molecules like SO2to violate the octetrule in order to be consistent with the experimentalobservation of equal bond lengths.

The octet on S expands.

S is in group 6, so up to 6 valence electrons can beinvolved in the bonding.

This is what occurs here. Note that there are normal octet structures one can write for SO2,

but these structure do not account for the experimental fact thatboth S-O bond lengths are identical.

The expansion occurs for third and higher period elements.

Other examples include SF6, SF4, PCl5.38

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Alert! Issue Expanded OctetPatch.

We Now Have Lewis Structures v1.2

For nonmetalsin the third and higher periodsit may not bepossible to satisfy the octet rule.

In these cases, we are permitted to expand the octet.

We are allowed to expand the octet in these elementsbecause, allegedly, dAOin the same valence shell (nquantum number) are not so high in energyfor theseelements and can thus participate in bonding.

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Al t! I R P t h

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O

O O

O

O O

O

O O

A

B

C

O

O O

A

B

C

Alert! Issue ResonancePatch.


The experimental fact was the bond lengths in O3are identical.

Hence the correct description of O2is not given by any one ofthe two Lewis structures individually but by the superposition of

the two, called a resonance hybrid.

The BO is 1 for each bond. 40

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Resonance Structures

They have the SAMErelative placement

of atoms BUT different locations ofbonding and lone electron pairs.

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Benzene C6H6

H

H

H

H

H

H

H

H

H

H

H

H

Experimentally, the bonds in benzene are all of equal length,

between a single and double bond.

The structure of benzene can also be written as......

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Benzene C6H6

Experimentally, the bonds in benzene are all of

equal length, between a single and double bond.

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Criteria for Choosing the More Important Resonance

structure:

Smaller formal charges are preferred over larger ones

(eg, 0 is preferred over -1 or +1)

The same nonzero formal charges on adjacent atoms are

not preferred

A more negative formal charge should reside on a moreelectronegative atom

Predicting Structures/shapes of molecules

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Predicting Structures/shapes of molecules

Lewis structures determine the bond pairs and lone

pairs of electrons among the atoms.

Combining this information with VSEPR(Valence Shell

Electron Pair Repulsiontheory) enables us to predict

the shapes of the molecules.

The basis of VSEPR is that the repulsions between

electrons in bonds and lone pairsdetermines the

overall shape of a molecule

VSEPR assumes that core electrons make no

significant impact on the shape of a molecule, so can

be ignored.45CM1502 Sem2-2013-14

VSEPR R l

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VSEPR Rules1. Draw Lewis structures to determine the bonding

around atoms.

2. Assign an electron group arrangement:- To do this we count the number of atoms

around the central atom and add it to the

number of lone pairs it possesses.

3. Predict the overall geometry around the central atom.

O in water has two hydrogen atomsbonded to it, and two lone pairs. So

there is 2+2=4 different directionselectrons are more localized in.

S in SO2has two oxygen atoms bondedto it, and a single lone pair. So there is2+1=3 different directions electrons are

more localized in.46

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VSEPR Rules

4. The final predicted structure is governed by the idea

that electrons in lp repel electrons found in other lpthe most,and electrons found in bonds repel other

electrons found in bonds the least,with lp electrons

repelling bonding electrons intermediatebetween the

above two.

5. Draw and name the molecular shape by counting

bonding groups and nonbonding groups separately.

6. Predict the bond angle.

movie (IVLE /workbin/videos/VSEPR theory47

VSEPR Structure Predictions

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VSEPR Structure Predictions

Electrons in 2 directions





Geometry about

central atom

Possible structures

of the molecule

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AXmEn Geometry structure examplesAX2 lilnear linear CS2,HCN, BeF2

AX3 Trigonal planar Trigonal planar SO3BF3NO3-CO32-

AX2E Trigonal planar V shaped SO2O3PbCl2SnBr2AX4 Tetrahedral Tetrahedral CH4SiCl4So4

2-ClO4-

AX3E Tetrahedral Trigonal pyramidal NH3PF3ClO3-H3O

+

AX2E2 Tetrahedral Bent H2O OF2SCl2

AX5 TBP TBP PF5AsF5SOF4

AX4E TBP Seesaw SF4XeO2F2IF4+IO2F2-

AX3E2 TBP Tshaped ClF3BrF3

AX2E3 TBP linear XeF2I3- IF2

-

AX6 Octahedral Octahedral SF6IOF5

AX5E Octahedral Squarepyramidal BrF5TeF5- XeOF4

AX4E2 Octahedral Square planar XeF4ICl4-

A - central atom, X - surrounding atom, E -lone pairs, m, n - integers 49CM1502 Sem2-2013-14

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Molecular shapes with more than one central atom

ethane

CH3CH3

ethanol

CH3CH2OH

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Conclusions

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Conclusions

VSEPR Each group of electrons around a

central atom remains as far away from the othersas possible.

5 common geometric shapesresult when 2, 3, 4,

5 or 6 electron groups surround a central atom.

Lone pairs and double bonds exert greater

repulsions.

Bond polarity and molecular shape determine

molecular polarity.

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