prentice hall © 2003chapter 8 chapter 8 basic concepts of chemical bonding chemistry the central...

Prentice Hall © 2003 Chapter 8

Chapter 8Chapter 8 Basic Concepts of Chemical Basic Concepts of Chemical

BondingBonding

CHEMISTRY The Central Science

9th Edition

David P. White


• 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.

Chemical Bonds, Lewis Chemical Bonds, Lewis Symbols, and the Octet Symbols, and the Octet

RuleRule


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 bonding are indicated by unpaired dots.

• These symbols are called Lewis symbols.• We generally place the electrons one four sides of a

square around the element symbol.


RuleRule


Lewis Symbols


RuleRule


The Octet Rule• All noble gases except He have an s2p6 configuration. • 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.


RuleRule


Consider the reaction between sodium and chlorine:Na(s) + ½Cl2(g) NaCl(s) Hºf = -410.9 kJ

Ionic BondingIonic Bonding


• The reaction is violently exothermic.• We infer that the NaCl is more stable than its constituent

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.




Energetics of Ionic Bond Formation• The formation of Na+(g) and Cl(g) from Na(g) and Cl(g)

is endothermic.• Why is the formation of NaCl(s) exothermic?• The reaction NaCl(s) Na+(g) + Cl(g) is endothermic

(H = +788 kJ/mol).• The formation of a crystal lattice from the ions in the gas

phase is exothermic:

Na+(g) + Cl(g) NaCl(s) H = 788 kJ/mol



Energetics of Ionic Bond Formation• 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 J·m/C2), Q1 and Q2 are the charges on the ions, and d is the distance between ions.

dQQ

El21



Energetics of Ionic Bond Formation• Lattice energy increases as

• The charges on the ions increase• The distance between the ions decreases.


Ionic BondingIonic BondingElectron Configurations of Ions of the

Representative Elements• These are derived from the electron configuration of

elements with the required number of electrons added or removed from the most accessible orbital.

• Electron configurations can predict stable ion formation:• Mg: [Ne]3s2

• Mg+: [Ne]3s1 not stable

• Mg2+: [Ne] stable

• Cl: [Ne]3s23p5

• Cl: [Ne]3s23p6 = [Ar] stable



Transition Metal Ions• Lattice energies compensate for the loss of up to three

electrons.• In general, electrons are removed from orbitals in order

of decreasing n (i.e. electrons are removed from 4s before the 3d).

Polyatomic Ions• Polyatomic ions are formed when there is an overall

charge on a compound containing covalent bonds. E.g. SO4

2, NO3.


Covalent BondingCovalent 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 connecting the two H nuclei.



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


Bond Polarity and Bond Polarity and ElectronegativityElectronegativity

• 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 atoms 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.

Bond Polarity and ElectronegativityBond Polarity and Electronegativity

Electronegativity



Electronegativity and Bond Polarity• Difference in electronegativity is a gauge of bond

polarity:• electronegativity differences around 0 result in non-polar

covalent bonds (equal or almost equal sharing of electrons);• electronegativity differences around 2 result in polar covalent

bonds (unequal sharing of electrons);• electronegativity differences around 3 result in ionic bonds

(transfer of electrons).



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


Bond Polarity and Bond Polarity and ElectronegativityElectronegativityBond Types and Nomenclature

• In general, the least electronegative element is named first.

• The name of the more electronegative element ends in –ide.

• Ionic compounds are named according to their ions, including the charge on the cation if it is variable.

• Molecular compounds are named with prefixes.


Bond Polarity and Bond Polarity and ElectronegativityElectronegativityBond Types and Nomenclature

Ionic Molecular

MgH2 Magnesium hydride H2S Hydrogen sulfide

FeF2 Iron(II) fluoride OF2 Oxygen difluoride

Mn2O3 Manganese(III) oxide

Cl2O3 Dichlorine trioxide


Drawing Lewis StructuresDrawing 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.



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.



Formal Charge• To calculate formal charge:

• All nonbonding electrons are assigned to the atom on which 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 3 from 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 3

from the triple bond. There are 5 electrons from the Lewis structure.

• Formal charge = 5 - 5 = 0.• We write:

C N

C N



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).

O

OO



Resonance Structures



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 structure has two identical bonds of intermediate character.

• Common examples: O3, NO3-, SO4

2-, NO2, and benzene.

O

OO

O

OO



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.



Resonance in Benzene• We write resonance structures for benzene in which

there are single bonds between each pair of C atoms and the 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 Exceptions to the Octet RuleRule

• 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 Groups 1A, 2A, and 3A.

• 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 extra electron density.


Strengths of Covalent Strengths of Covalent BondsBonds

• The energy required to break a covalent bond is called the bond dissociation enthalpy, D. That is, for the Cl2 molecule, D(Cl-Cl) is given by H for 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 the

atomization reaction:D(C-H) = ¼H = ¼(1660 kJ) = 415 kJ.

• Bond enthalpies are always positive.



Bond Enthalpies and the Enthalpies of Reactions

• 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.




• Mathematically, if Hrxn 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 = ?

formed bondsbroken bonds DDHrxn

Strengths of Covalent BondsStrengths of Covalent Bonds




• 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 that the bonds formed are stronger than the bonds broken.

• The above result is consistent with Hess’s law.

kJ 104

Cl-HCl-CCl-ClH-C

DDDDHrxn



Bond Enthalpy and Bond Length• We know that multiple bonds are shorter than single

bonds.• 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.

Strengths of Covalent BondsStrengths of Covalent Bonds


End of Chapter 8:End of Chapter 8:Basic Concepts of Chemical Basic Concepts of Chemical

BondingBonding

prentice hall © 2003chapter 8 chapter 8 basic concepts of chemical bonding chemistry the central...

Documents

na ions

octet of electrons

ionic lattice

gaseous ions

ionic bond results

ionic bondingnacl

reaction nacls na g

sharing electrons