Bonding 1 2012

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Chapter 8Basic Concepts of Chemical

Bonding

CHEMISTRYThe Central Science

9th Edition

Cindy FuhrerHuntley High School

Chemical bond: attractive force holding two or more

atoms together.

Covalent bond results from sharing electrons between

the atoms. Usually found between nonmetals.

Ionic bond results from the transfer of electrons from a

metal to a nonmetal. Metallic bond: attractive force holding pure metals

together. (electron sea)

Chemical Bonds, Lewis

Symbols, and the OctetRule

Lewis Symbols

As a pictorial understanding of where the

electrons are in an atom, we represent the

electrons as dots around the symbol for the

element.

The number of electrons available for bondingare indicated by unpaired dots.

These symbols are called Lewis symbols.

We generally place the electrons on four sides

of a square around the element symbol.

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The Octet Rule

All noble gases except He have an s2p6configuration.

Octet rule: atoms tend to gain, lose, or share

electrons until they are surrounded by 8

valence electrons (4 electron pairs).

Caution: there are many exceptions to the

octet rule.

Consider the reaction between sodium and chlorine:

Na(s) + Cl2(g) NaCl(s) Hf= -410.9 kJ

Ionic Bonding

The reaction is violently exothermic.

We infer that the NaCl is more stable than its

elements. Why?

Na has lost an electron to become Na+ and

chlorine has gained the electron to become Cl

.Note: Na+ has an Ne electron configuration and

Cl has an Ar configuration.

That is, both Na+ and Cl have an octet of

electrons surrounding the central ion.

NaCl forms a very regular structure in which

each Na+ ion is surrounded by 6 Cl ions.

Similarly, each Cl ion is surrounded by six

Na+ ions.

There is a regular arrangement of Na+ and Cl

in 3D.

Note that the ions are packed as closely as

possible.

Note that it is not easy to find a molecular

formula to describe the ionic lattice.

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Lattice energy: the energy required to completely

separate an ionic solid into its gaseous ions.

Lattice energy depends on the charges on the ions

and the sizes of the ions:

is a constant (8.99 x 10 9 Jm/C2), Q1 and Q2 arethe charges on the ions, and dis the distance

between ions.

Lattice energy increases asThe charges on the ions increase

The distance between the ions decreases.

d

QQEl

21=

9.3

Lattice energy (E)increases as Q

increases and/oras r decreases.

cmpd lattice energyMgF2MgO

LiF

LiCl

29573938

1036

853

Q= +2,-1

Q= +2,-2

r F < r Cl

Electrostatic (Lattice) Energy

E =kQ+Q-r

Q+ is the charge on the cationQ- is the charge on the anion

r is the distance between the ions

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Examples Which in each pair

would have the greatest latticeenergy?

1. LiF or LiCl

2. NaCl or MgCl23. KBr or KI

4. MgCl2 or MgO

5. CaO or NaF

Examples Which in each pair

would have the greatest latticeenergy?

1. LiF or LiCl

2. NaCl or MgCl23. KBr or KI

4. MgCl2 or MgO

5. CaO or NaF

1. Vaporize the metal (enthalpy ofvaporization)

Na (s) Na (g)

2. Break diatomic nonmetal molecules (ifapplicable) (bond enthalpy)

Cl2 (g) Cl-

3. Remove electron(s) from metal (ionizationenergy)

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

Born-Haber CycleApplication of Hesss Law

4. Add electron(s) to nonmetal (electronaffinity)

Cl (g) + e- Cl- (g)

5. Put ions together to form compound (latticeenergy)

Na+ + Cl- NaCl (s)

Born-Haber CycleApplication of Hesss Law

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Overall Reaction:

Na (s) + Cl2 (g) NaCl (s)

This is useful because allquantities are directly measurableexcept lattice energy. The Born-Haber cycle can be used tocalculate lattice energy from theother values.

5

4

3

2

1

Step

HStep

1 Enthalpy of Vaporization endothermic

2 Bond Enthalpy

endothermic

3 Ionization Energy endothermic

4 Electron Affinity Exothermic

5 Lattice Energy Exothermic (highly)

9.3

Born-Haber Cycle for Determining Lattice Energy, page

Hoverall = H1 + H2 + H3 + H4 + H5o ooooo

Heat of Vap

Bond Enthalpy

Ionization Energy

Electron Affinity

Lattice Energy

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Covalent Bonding

When two similar atoms bond, none of them wants to

lose or gain an electron to form an octet.

When similar atoms bond, they share pairs of electrons to

each obtain an octet.

Each pair of shared electrons constitutes one chemical

bond.

Example: H + H H2 has electrons on a line connectingthe two H nuclei.

Covalent Bonding

Lewis Structures

Covalent bonds can be represented by the Lewis symbols

of the elements:

In Lewis structures, each pair of electrons in a bond is

represented by a single line:

Cl + Cl Cl Cl

Cl Cl H FH O

H

H N H

HCH

H

H

H

Multiple Bonds

It is possible for more than one pair of electrons to be

shared between two atoms (multiple bonds):

One shared pair of electrons = single bond (e.g. H2);

Two shared pairs of electrons = double bond (e.g. O2);

Three shared pairs of electrons = triple bond (e.g. N2).

Generally, bond distances decrease as we move from

single through double to triple bonds.

H H O O N N

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Bond Polarity and

Electronegativity In a covalent bond, electrons are shared.

Sharing of electrons to form a covalent bond does not

imply equal sharing of those electrons.

There are some covalent bonds in which the electrons are

located closer to one atom than the other.

Unequal sharing of electrons results in polar bonds.

Electronegativity

Electronegativity: The ability of one atom in a molecule

to attract electrons to itself.

Pauling set electronegativities on a scale from 0.7 (Cs) to

4.0 (F).

Electronegativity increases

across a period and

up a group.

Electronegativity and Bond Polarity

Difference in electronegativity is a gauge of bond

polarity or location of the atoms on the periodic table:

Nonpolar Covalent Bond - equal or almost equal sharing of

electrons, electronegativity difference of 0 -0.3 or both

nonmetals

Polar Covalent Bond- unequal sharing of electrons,electronegativity differences 0.4-1.6

Ionic Bond - transfer of electrons, electronegativity difference

greater than 1.7, metal to nonmetal especially group 1 or 2 to

16 or 17

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Electronegativity and Bond Polarity

There is no sharp distinction between bonding types.

The positive end (or pole) in a polar bond is represented

+ and the negative pole -.

Dipole Moments

Consider HF:

The difference in electronegativity leads to a polar bond.

There is more electron density on F than on H.

Since there are two different ends of the molecule, we call HF

a dipole.

Dipole moment, , is the magnitude of the dipole:

where Q is the magnitude of the charges.

Dipole moments are measured in debyes, D.

Qr=

Drawing Lewis Structures

1. Add the valence electrons.

2. Write symbols for the atoms and show which atoms are

connected to which.

3. Complete the octet for the central atom, then complete

the octets of the other atoms.

4. Place leftover electrons on the central atom.

5. If there are not enough electrons to give the central atom

an octet, try multiple bonds.

Examples Draw Lewis Structuresfor each:

1. H2O2. CO23. NCl34. SO25. SO3

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Formal Charge

It is possible to draw more than one Lewis structure with

the octet rule obeyed for all the atoms.

To determine which structure is most reasonable, we use

formal charge.

Formal charge is the charge on an atom that it would

have if all the atoms had the same electronegativity.

To calculate formal charge:

All nonbonding electrons are assigned to the atom onwhich they are found.

Half the bonding electrons are assigned to each atom

in a bond.

Formal charge is:

valence electrons - number of bonds - lone pair electrons

Formal Charge

Consider:

For C:

There are 4 valence electrons (from periodic table).

In the Lewis structure there are 2 nonbonding electrons and 3from the triple bond. There are 5 electrons from the Lewis

structure.

Formal charge: 4 - 5 = -1.

C N

Formal Charge

Consider:

For N: There are 5 valence electrons.

In the Lewis structure there are 2 nonbonding electrons and 3from the triple bond. There are 5 electrons from the Lewis

structure.

Formal charge = 5 - 5 = 0.

We write:

C N

C N

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Examples Determine the formal charge on each

atom for each:

1.H2O

2.CO2

3.NCl3

4.SO2

5.SO3

Formal Charge The most stable structure has:

the lowest formal charge on each atom, the most negative formal charge on the most electronegative

atoms.

Resonance Structures Some molecules are not well described by Lewis

Structures.

Typically, structures with multiple bonds can have

similar structures with the multiple bonds between

different pairs of atoms

Resonance Structures Example: experimentally, ozone has two identical bonds

whereas the Lewis Structure requires one single (longer)

and one double bond (shorter).

OO O

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Resonance Structures Resonance structures are attempts to represent a real

structure that is a mix between several extreme

possibilities.

Resonance Structures Example: in ozone the extreme possibilities have one

double and one single bond. The resonance structurehas two identical bonds of intermediate character.

Common examples: O3, NO3-, SO4

2-, NO2, and benzene.

O

OO

O

OO

Examples Draw Lewis Structures for

each, and include all relevantresonance structures:

1. NO3-

2. CO32-

3. NO2

Resonance in Benzene Benzene consists of 6 carbon atoms in a hexagon. Each

C atom is attached to two other C atoms and one

hydrogen atom.

There are alternating double and single bonds between

the C atoms.

Experimentally, the C-C bonds in benzene are all the

same length.

Experimentally, benzene is planar.

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We write resonance structures for benzene in which

there are single bonds between each pair of C atoms andthe 6 additional electrons are delocalized over the entire

ring:

Benzene belongs to a category of organic molecules

called aromatic compounds (due to their odor).

or

Exceptions to the Octet

Rule

There are three classes of exceptions to the octet rule:

Molecules with an odd number of electrons;

Molecules in which one atom has less than an octet;

Molecules in which one atom has more than an octet.

Odd Number of Electrons

Few examples. Generally molecules such as ClO2, NO,and NO2 have an odd number of electrons.

N O N O

Less than an Octet

Relatively rare.

Molecules with less than an octet are typical for

compounds of Group 13.

Most typical example is BF3.

Formal charges indicate that the Lewis structure with an

incomplete octet is more important than the ones with

double bonds.

More than an Octet

This is the largest class of exceptions.

Atoms from the 3rd period onwards can accommodate

more than an octet.

Beyond the third period, the d-orbitals are low enough in

energy to participate in bonding and accept the extraelectron density.

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Examples Draw Lewis Structuresfor Each

1. PCl52. SF63. BeCl24. BF3

5. XeF4

Strengths of Covalent

Bonds

The energy required to break a covalent bond is called

the bond dissociation enthalpy,D. That is, for the Cl2molecule,D(Cl-Cl) is given by Hfor the reaction:

Cl2(g) 2Cl(g). When more than one bond is broken:

CH4(g) C(g) + 4H(g) H= 1660 kJ

the bond enthalpy is a fraction of H for theatomization reaction:

D(C-H) = H= (1660 kJ) = 415 kJ. Bond enthalpies can either be positive or negative.

Bond Enthalpies and the Enthalpies ofReactions

We can use bond enthalpies to calculate the enthalpy for

a chemical reaction.

We recognize that in any chemical reaction bonds need

to be broken and then new bonds get formed. The enthalpy of the reaction is given by the sum of bond

enthalpies for bonds broken less the sum of bond

enthalpies for bonds formed.

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Mathematically, ifHrxn is the enthalpy for a reaction,then

We illustrate the concept with the reaction between

methane, CH4, and chlorine:

CH4(g) + Cl2(g) CH3Cl(g) + HCl(g) Hrxn = ?

( ) ( ) = formedbondsbrokenbonds DDHrxn

In this reaction one C-H bond and one Cl-Cl bond gets

broken while one C-Cl bond and one H-Cl bond gets

formed.

The overall reaction is exothermic which means than the

bonds formed are stronger than the bonds broken.

The above result is consistent with Hesss law.

( ) ( )[ ] ( ) ( )[ ]{ }

kJ104

Cl-HCl-CCl-ClH-C

=

++= DDDDHrxn

Example Use bond energies fromtable 8.4 to calculate H for the

following reaction:

C C

C

H

H

H

H H

H + C C C

H

H

H

H

H

H

H Cl

H

Cl

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Bond Enthalpy and Bond Length

We know that multiple bonds are shorter than singlebonds.

We can show that multiple bonds are stronger than

single bonds.

As the number of bonds between atoms increases, the

atoms are held closer and more tightly together.

NOTE: A double bond between two atoms is not twice

as strong as a single bond between the same two atoms.