- Ionic bonds form when a metal reacts with a nonmetal, resulting in the transfer of electrons from the metal to the nonmetal. The ions achieve noble gas electron configurations. - Covalent bonds form when two nonmetals share electrons between their nuclei. Polar covalent bonds result in unequal sharing of electrons between the atoms. - Electronegativity describes an atom's ability to attract shared electrons. It increases left to right and decreases top to bottom on the periodic table. Unequal electronegativity between bonded atoms results in bond polarity.

1. Bonding: GeneralBonding: General
ConceptsConcepts
Advanced Chemistry Chapter 8Advanced Chemistry Chapter 8

2. Types of ChemicalTypes of Chemical
BondsBonds

3. Ionic BondsIonic Bonds
Ionic Bonds are formed when an atomIonic Bonds are formed when an atom
that loses electrons relatively easilythat loses electrons relatively easily
reacts with an atom that has a highreacts with an atom that has a high
attraction for electrons.attraction for electrons.
Ionic Compounds results when a metalIonic Compounds results when a metal
bonds with a nonmetal.bonds with a nonmetal.

4. Bond EnergyBond Energy
Bond energy is the energy required to break aBond energy is the energy required to break a
bond.bond.
The energy of interaction between a pair of ionsThe energy of interaction between a pair of ions
can be calculated using Coulomb’s lawcan be calculated using Coulomb’s law
r = the distance between the ions in nm.r = the distance between the ions in nm.
QQ11 and Qand Q22 are the numerical ion charges.are the numerical ion charges.
E is in joulesE is in joules
E = (2.31ξ10−19
ϑγνµ )
Θ1Θ2
ρ





5. Bond EnergyBond Energy
When the calculated energy betweenWhen the calculated energy between
ions is negative, that indicates anions is negative, that indicates an
attractive force.attractive force.
A positive energy is a repulsive energy.A positive energy is a repulsive energy.
The distance where the energy isThe distance where the energy is
minimal is called the bond length.minimal is called the bond length.

6. Covalent BondsCovalent Bonds
Covalent bonds form betweenCovalent bonds form between
molecules in which electrons are sharedmolecules in which electrons are shared
by nuclei.by nuclei.
The bonding electrons are typicallyThe bonding electrons are typically
positioned between the two positivelypositioned between the two positively
charged nuclei.charged nuclei.

7. Polar Covalent BondsPolar Covalent Bonds
Polar covalent bonds are an intermediatePolar covalent bonds are an intermediate
case in which the electrons are notcase in which the electrons are not
completely transferred but form unequalcompletely transferred but form unequal
sharing.sharing.
A δA δ--
or δor δ++
is used to show a fractional or partialis used to show a fractional or partial
charge on a molecule with unequal sharing.charge on a molecule with unequal sharing.
This is called a dipole.This is called a dipole.

8. ElectronegativityElectronegativity

9. ElectronegativityElectronegativity
Electronegativity is the ability of an atom in aElectronegativity is the ability of an atom in a
molecule to attract shared electrons to itself.molecule to attract shared electrons to itself.
(electron love)(electron love)
Relative electronegativities are determined byRelative electronegativities are determined by
comparing the measured bond energy with thecomparing the measured bond energy with the
“expected” bond energy.“expected” bond energy.
Measured in Paulings. After Linus Pauling theMeasured in Paulings. After Linus Pauling the
American scientist who won the Nobel PrizesAmerican scientist who won the Nobel Prizes
for both chemistry and peace.for both chemistry and peace.

10. ElectronegativityElectronegativity
Expected H-X bond energy=Expected H-X bond energy=
H − Η βονδ ενεργψ+ Ξ − Ξ βονδ ενεργψ
2

11. ElectronegativityElectronegativity
Electronegativity values generally increaseElectronegativity values generally increase
going left to right across the periodic tablegoing left to right across the periodic table
and decrease going top to bottom.and decrease going top to bottom.

12. Electronegativity andElectronegativity and
Bond typeBond type

13. Bond Polarity and DipoleBond Polarity and Dipole

14. Dipoles and DipoleDipoles and Dipole
MomentsMoments
A molecule that has a center of positiveA molecule that has a center of positive
charge and a center of negative chargecharge and a center of negative charge
is said to beis said to be dipolardipolar or to have aor to have a dipoledipole
moment.moment.
An arrow is used to show this dipoleAn arrow is used to show this dipole
moment by pointing to the negativemoment by pointing to the negative
charge and the tail at the positivecharge and the tail at the positive
charge.charge.

15. Dipoles and DipoleDipoles and Dipole
MomentsMoments
Electrostatic potentialElectrostatic potential
diagram showsdiagram shows
variation in charge.variation in charge.
Red is the mostRed is the most
electron rich regionelectron rich region
and blue is the mostand blue is the most
electron poor region.electron poor region.

16. Dipoles and DipoleDipoles and Dipole
MomentsMoments

17. Dipoles and DipoleDipoles and Dipole
MomentsMoments

18. Dipoles and DipoleDipoles and Dipole
MomentsMoments

19. Dipoles and DipoleDipoles and Dipole
MomentsMoments
Dipole moments are when opposingDipole moments are when opposing
bond polarities don’t cancel out.bond polarities don’t cancel out.

20. Dipoles and DipoleDipoles and Dipole
MomentsMoments

21. Example ProblemsExample Problems
For each of the following molecules,For each of the following molecules,
show the direction of the bondshow the direction of the bond
polarities and indicate which ones havepolarities and indicate which ones have
a dipole moment: HCl, Cla dipole moment: HCl, Cl22, SO, SO33, CH, CH44, H, H22SS

22. HClHCl

23. ClCl22

24. SOSO33

25. CHCH44

26. HH22SS

27. Ions: ElectronIons: Electron
Configurations and SizesConfigurations and Sizes

28. Electron ConfigurationsElectron Configurations
of Compoundsof Compounds
When two nonmetals react to form aWhen two nonmetals react to form a
covalent bond, they share electrons in a waycovalent bond, they share electrons in a way
that completes the valence electronthat completes the valence electron
configurations of both atoms. That is, bothconfigurations of both atoms. That is, both
nonmetals attain noble gas electronnonmetals attain noble gas electron
configurations.configurations.

29. Electron ConfigurationsElectron Configurations
of Compoundsof Compounds
When a nonmetal and a representative-groupWhen a nonmetal and a representative-group
metal react to form a binary ionicmetal react to form a binary ionic
compounds, the ions form so that the valencecompounds, the ions form so that the valence
electron configuration of the nonmetalelectron configuration of the nonmetal
achieves the electron configuration of theachieves the electron configuration of the
next noble gas atom and the valence orbitalsnext noble gas atom and the valence orbitals
of the metal are emptied. In this way bothof the metal are emptied. In this way both
ions achieve noble gas electronions achieve noble gas electron
configurations.configurations.

30. Predicting IonicPredicting Ionic
FormulasFormulas
To predict the formula of the ionicTo predict the formula of the ionic
compound, we simply recognize that thecompound, we simply recognize that the
chemical compounds are always electricallychemical compounds are always electrically
neutral. They have the same quantities ofneutral. They have the same quantities of
positive and negative charges.positive and negative charges.

31. Sizes of IonsSizes of Ions
Size of an ion generally follows the same trendSize of an ion generally follows the same trend
as atomic radius. The big exception to thisas atomic radius. The big exception to this
trend is where the metals become nonmetalstrend is where the metals become nonmetals
and the ions switch charge.and the ions switch charge.

32. Sizes of IonsSizes of Ions
A positive ion is formed by removing one orA positive ion is formed by removing one or
more electrons from a neutral atom, themore electrons from a neutral atom, the
resulting cation is smaller than the neutralresulting cation is smaller than the neutral
atom.atom.
Less electrons allow for less repulsions andLess electrons allow for less repulsions and
the ion gets smaller.the ion gets smaller.

33. Sizes of IonsSizes of Ions
An addition of electrons to a neutral atomAn addition of electrons to a neutral atom
produces an anion that is significantly largerproduces an anion that is significantly larger
than the neutral atom.than the neutral atom.
An addition of an electron causes additionalAn addition of an electron causes additional
repulsions around the atom and therefore itsrepulsions around the atom and therefore its
size increases.size increases.

34. Energy Effects in BinaryEnergy Effects in Binary
Ionic CompoundsIonic Compounds

35. Lattice EnergyLattice Energy
Lattice energy is the change in energy that takesLattice energy is the change in energy that takes
place when separated gaseous ions areplace when separated gaseous ions are
packed together to form an ionic solid.packed together to form an ionic solid.
The lattice energy is often defined as theThe lattice energy is often defined as the
energy released when an ionic solid formsenergy released when an ionic solid forms
from its ions.from its ions.
Lattice energy has a negative sign to showLattice energy has a negative sign to show
that the energy is released.that the energy is released.

36. Lattice Energy ExampleLattice Energy Example
Estimate the enthalpy of lithium fluoride andEstimate the enthalpy of lithium fluoride and
the changes of energy and lattice energythe changes of energy and lattice energy
during formation:during formation:
1.1. Break down LiF into its standard stateBreak down LiF into its standard state
elements (use formation reaction):elements (use formation reaction):
LiLi(s)(s) + ½F+ ½F2(g)2(g)  LiFLiF(s)(s)
Li+
(g) + F-
(g)  LiF(s)

37. Lattice Energy ExampleLattice Energy Example
LiLi(s)(s) + ½F+ ½F2(g)2(g)  LiFLiF(s)(s) LiLi++
(g)(g) + F+ F--
(g)(g)  LiFLiF(s)(s)
Use sublimation and evaporation reactions to getUse sublimation and evaporation reactions to get
reactants into gas form (since lattice energyreactants into gas form (since lattice energy
depends on gaseous state). Find the enthalpies todepends on gaseous state). Find the enthalpies to
these reactions:these reactions:
LiLi(s)(s)  LiLi(g)(g) 161 kJ/mol161 kJ/mol
LiLi(g)(g) + ½F+ ½F2(g)2(g)  LiFLiF(s)(s)

38. Lattice Energy ExampleLattice Energy Example
LiLi(g)(g) + ½F+ ½F2(g)2(g)  LiFLiF(s)(s) LiLi++
(g)(g) + F+ F--
(g)(g)  LiFLiF(s)(s)
Ionize cation to form ions for bonding. UseIonize cation to form ions for bonding. Use
Ionization energy for the enthalpy of the reaction.Ionization energy for the enthalpy of the reaction.
LiLi(g)(g)  LiLi++
(g)(g) + e+ e--
Ionization energy: 520 kJ/molIonization energy: 520 kJ/mol
LiLi++
(g)(g) + ½F+ ½F2(g)2(g)  LiFLiF(s)(s)

39. Lattice Energy ExampleLattice Energy Example
LiLi++
(g)(g) + ½F+ ½F2(g)2(g)  LiFLiF(s)(s) LiLi++
(g)(g) + F+ F--
(g)(g)  LiFLiF(s)(s)
Dissociate diatomic gas to individual atoms:Dissociate diatomic gas to individual atoms:
½F½F2(g)2(g)  FF(g)(g) ½ Bond dissociation energy of F-F½ Bond dissociation energy of F-F
= 154 kJ/ 2 = 77 kJ/mol= 154 kJ/ 2 = 77 kJ/mol
LiLi++
(g)(g) + F+ F(g)(g)  LiFLiF(s)(s)

40. Lattice Energy ExampleLattice Energy Example
LiLi++
(g)(g) + F+ F(g)(g)  LiFLiF(s)(s) LiLi++
(g)(g) + F+ F--
(g)(g)  LiFLiF(s)(s)
Electron addition to fluorine is the electron affinityElectron addition to fluorine is the electron affinity
of fluorine:of fluorine:
FF(g)(g) + e+ e--
 FF--
(g)(g) -328 kJ/mol-328 kJ/mol
LiLi++
(g)(g) + F+ F--
(g)(g)  LiFLiF(s)(s)

41. Lattice Energy ExampleLattice Energy Example
LiLi++
(g)(g) + F+ F--
(g)(g)  LiFLiF(s)(s) LiLi++
(g)(g) + F+ F--
(g)(g)  LiFLiF(s)(s)
Formation of solid lithium fluoride from theFormation of solid lithium fluoride from the
gaseous ions corresponds to its lattice energy:gaseous ions corresponds to its lattice energy:
LiLi++
(g)(g) + F+ F--
(g)(g)  LiFLiF(s)(s) -1047 kJ/mol-1047 kJ/mol

42. Lattice Energy ExampleLattice Energy Example
The sum of these five processes yields the overallThe sum of these five processes yields the overall
reaction and the sum of the individual energy changesreaction and the sum of the individual energy changes
gives the overall energy change or lattice energy:gives the overall energy change or lattice energy:
LiLi(s)(s)  LiLi(g)(g)
LiLi(g)(g)  LiLi++
(g)(g) + e+ e--
½F½F2(g)2(g)  FF(g)(g)
FF(g)(g) + e+ e--
 FF--
(g)(g)
LiLi++
(g)(g) + F+ F--
(g)(g)  LiFLiF(s)(s)
161 kJ161 kJ
520 kJ520 kJ
77 kJ77 kJ
-328 kJ-328 kJ
-1047 kJ-1047 kJ
Total = -617 kJ/molTotal = -617 kJ/mol

43. Lattice EnergyLattice Energy

44. Lattice EnergyLattice Energy
Lattice energy can be calculated with at form ofLattice energy can be calculated with at form of
Coulomb’s law:Coulomb’s law:
Q is the charges on the ions and r is theQ is the charges on the ions and r is the
shortest distance between the centers of theshortest distance between the centers of the
cations and anions. k is a constant thatcations and anions. k is a constant that
depends on the structure of the solid and thedepends on the structure of the solid and the
electron configurations of the ions.electron configurations of the ions.
LatticeEnergy = κ
Θ1Θ2
ρ





45. Partial Ionic CharacterPartial Ionic Character
of Covalent Bondsof Covalent Bonds

46. Bond CharacterBond Character
Calculations of ionic character:Calculations of ionic character:
Even compounds with the maximum possibleEven compounds with the maximum possible
electronegativity differences are not 100%electronegativity differences are not 100%
ionic in the gas phase. Therefore theionic in the gas phase. Therefore the
operational definition of ionic is anyoperational definition of ionic is any
compound that conducts an electric currentcompound that conducts an electric current
when melted will be classified as ionic.when melted will be classified as ionic.
Percent ionic character of a bond =
διπολε µοµεντ οφξ − ψ
διπολε µοµεντ οφξ+
ψψ



 ξ100%

47. Bond CharacterBond Character

48. The Covalent ChemicalThe Covalent Chemical
BondBond

49. Chemical Bond ModelChemical Bond Model
A chemical bond can be viewed as forces thatA chemical bond can be viewed as forces that
cause a group of atoms to behave as a unit.cause a group of atoms to behave as a unit.
Bonds result from the tendency of a systemBonds result from the tendency of a system
to seek its lowest possible energy.to seek its lowest possible energy.
Individual bonds act relatively independent.Individual bonds act relatively independent.

50. ExampleExample
It takes 1652 kJ of energy required to break theIt takes 1652 kJ of energy required to break the
bonds in 1 mole of methane.bonds in 1 mole of methane.
1652 kJ of energy is released when 1 mole of1652 kJ of energy is released when 1 mole of
methane is formed from gaseous atoms.methane is formed from gaseous atoms.
Therefore, 1 mole of methane in gas phase hasTherefore, 1 mole of methane in gas phase has
1652 kJ lower energy than the total of the1652 kJ lower energy than the total of the
individual atoms.individual atoms.
One mole of methane is held together with 1652One mole of methane is held together with 1652
kJ of energy.kJ of energy.
Each of the four C-H bonds contains 413 kJ ofEach of the four C-H bonds contains 413 kJ of
energy.energy.

51. ExampleExample
Each of the four C-H bonds contains 413 kJEach of the four C-H bonds contains 413 kJ
of energy.of energy.
CHCH33Cl contains 1578 kJ of energy:Cl contains 1578 kJ of energy:
1 mol of C-Cl bonds + 3 mol (C-H bonds)=1578 kJ1 mol of C-Cl bonds + 3 mol (C-H bonds)=1578 kJ
C-Cl bond energy + 3 (413 kJ/mol) = 1578 kJC-Cl bond energy + 3 (413 kJ/mol) = 1578 kJ
C-Cl bond energy = 339 kJ/molC-Cl bond energy = 339 kJ/mol

52. Properties of ModelsProperties of Models
A model doesn’t equal reality; they are usedA model doesn’t equal reality; they are used
to explain incomplete understanding of howto explain incomplete understanding of how
nature works.nature works.
Models are often oversimplified and areModels are often oversimplified and are
sometimes wrong.sometimes wrong.
Models over time tend to get overModels over time tend to get over
complicated due to “repairs”.complicated due to “repairs”.

53. Properties of ModelsProperties of Models
Remember that simple models often requireRemember that simple models often require
restrictive assumptions. Best way to userestrictive assumptions. Best way to use
models is to understand their strengths andmodels is to understand their strengths and
weaknesses.weaknesses.
We often learn more when models areWe often learn more when models are
incorrect than when they are right.incorrect than when they are right.
Cu and Cr.Cu and Cr.

54. Covalent Bond EnergiesCovalent Bond Energies
and Chemical Reactionsand Chemical Reactions

55. Bond EnergiesBond Energies
Bond energy averages are used for individualBond energy averages are used for individual
bond dissociation energies to givebond dissociation energies to give
approximate energies in a particular bond.approximate energies in a particular bond.
Bond energies vary due to several reasons:Bond energies vary due to several reasons:
multiple bonds, 4 C-H bonds in methanemultiple bonds, 4 C-H bonds in methane
different elements in the molecule, C-H bond indifferent elements in the molecule, C-H bond in
CC22HH66 or C-H bond in HCClor C-H bond in HCCl33

56. Bond Energy ExampleBond Energy Example
CHCH4(g)4(g)CHCH3(g)3(g) + H+ H(g)(g)
CHCH3(g)3(g)CHCH2(g)2(g) + H+ H(g)(g)
CHCH2(g)2(g)CHCH(g)(g) + H+ H(g)(g)
CHCH(g)(g)CC(g)(g) + H+ H(g)(g)
435 kJ435 kJ
453 kJ453 kJ
425 kJ425 kJ
339 kJ339 kJ
Total 1652 kJTotal 1652 kJ
Average 413 kJAverage 413 kJ

57. Bond Energy ExampleBond Energy Example
HCBrHCBr33
HCClHCCl33
HCFHCF33
CC22HH66
380 kJ380 kJ
380 kJ380 kJ
430 kJ430 kJ
410 kJ410 kJ

58. Average Bond EnergiesAverage Bond Energies

59. Bond EnergyBond Energy
A relationship also exists between theA relationship also exists between the
number of shared electron pairs.number of shared electron pairs.
single bond – 2 electronssingle bond – 2 electrons
double bond – 4 electronsdouble bond – 4 electrons
triple bond – 6 electronstriple bond – 6 electrons

60. Bond EnergyBond Energy
Bond energy values can be used to calculateBond energy values can be used to calculate
approximate energies for reactions.approximate energies for reactions.
Energy associated with bond breaking haveEnergy associated with bond breaking have
positive signspositive signs
Endothermic processEndothermic process
Energy associated with forming bonds releasesEnergy associated with forming bonds releases
energy and is negative.energy and is negative.
Exothermic processExothermic process

61. Bond EnergyBond Energy
A relationship exists between the number ofA relationship exists between the number of
shared electron pairs and the bond length.shared electron pairs and the bond length.
• As the number of electrons shared goes up theAs the number of electrons shared goes up the
bond length shortens.bond length shortens.

62. Bond EnergyBond Energy

63. Bond EnergyBond Energy
ΔH = sum of the energies required to break oldΔH = sum of the energies required to break old
bonds (positive signs) plus the sum of thebonds (positive signs) plus the sum of the
energies released in the formation of new bondsenergies released in the formation of new bonds
(negative signs).(negative signs).
• D represents bond energies per mole and alwaysD represents bond energies per mole and always
has positive signhas positive sign
• n is number of molesn is number of moles
∆H = Σn x D(bonds broken) − ∑n x D(bonds formed)

64. Bond Energy ExampleBond Energy Example
HH2(g)2(g) + F+ F2(g)2(g) 2HF2HF(g)(g)
1 H-H bond, F-F bond and 2 H-F bonds1 H-H bond, F-F bond and 2 H-F bonds
ΔH = DΔH = DH-HH-H + D+ DF-FF-F – 2D– 2DH-FH-F
• ΔH= (1mol x 432 kJ/mol) + (1mol x 154 kJ/mol)ΔH= (1mol x 432 kJ/mol) + (1mol x 154 kJ/mol)
– (2mol x 565 kJ/mol)– (2mol x 565 kJ/mol)
ΔH = -544 kJΔH = -544 kJ

65. The Localized ElectronThe Localized Electron
Bonding ModelBonding Model

66. Localized ElectronLocalized Electron
ModelModel
• The localized electron model assumes that aThe localized electron model assumes that a
molecule is composed of atoms that aremolecule is composed of atoms that are
bound together by sharing pairs of electronsbound together by sharing pairs of electrons
using the atomic orbitals of the bound atoms.using the atomic orbitals of the bound atoms.
• Electrons are assumed to be localized on aElectrons are assumed to be localized on a
particular atom individually or in the spaceparticular atom individually or in the space
between atoms.between atoms.

67. Localized ElectronLocalized Electron
ModelModel
• Pairs of electrons that are localized on anPairs of electrons that are localized on an
atom are called lone pairs.atom are called lone pairs.
• Pairs of electrons that are found in the spacePairs of electrons that are found in the space
between the atoms are called bonding pairsbetween the atoms are called bonding pairs

68. Localized ElectronLocalized Electron
ModelModel
Three parts of the LE Model:Three parts of the LE Model:
1.1. Description of the valence electron arrangementDescription of the valence electron arrangement
in the molecule using Lewis structures.in the molecule using Lewis structures.
2.2. Prediction of the geometry of the molecule usingPrediction of the geometry of the molecule using
VSEPR modelVSEPR model
3.3. Description of the type of atomic orbitals used byDescription of the type of atomic orbitals used by
the atoms to share electrons or hold lone pairs.the atoms to share electrons or hold lone pairs.

69. Lewis StructuresLewis Structures

70. Lewis StructuresLewis Structures
The Lewis structure of a molecule show how theThe Lewis structure of a molecule show how the
valence electrons are arranged among the atomsvalence electrons are arranged among the atoms
in the molecule.in the molecule.
• Named after G. N. LewisNamed after G. N. Lewis
• Rules are based on observations of thousands ofRules are based on observations of thousands of
molecules.molecules.
• Most important requirement for the formation ofMost important requirement for the formation of
a stable compound is that the atoms achieve noblea stable compound is that the atoms achieve noble
gas electron configurations.gas electron configurations.

71. Lewis StructuresLewis Structures
• Only the valence electrons are included.Only the valence electrons are included.
• The duet rule: diatomic molecules can findThe duet rule: diatomic molecules can find
stability in the sharing of two electrons.stability in the sharing of two electrons.
• The octet rule: since eight electrons areThe octet rule: since eight electrons are
required to fill these orbitals, these elementsrequired to fill these orbitals, these elements
typically are surrounded by eight electrons.typically are surrounded by eight electrons.

72. Lewis Structure StepsLewis Structure Steps
1.1. Sum the valence electrons from all the atoms.Sum the valence electrons from all the atoms.
Total valence electrons.Total valence electrons.
2.2. Use a pair of electrons to form a bond betweenUse a pair of electrons to form a bond between
each pair of bound atoms.each pair of bound atoms.
3.3. Arrange the remaining electrons to satisfy theArrange the remaining electrons to satisfy the
duet rule for hydrogen and the octet rule for theduet rule for hydrogen and the octet rule for the
others.others.
a)a) Terminal atoms first.Terminal atoms first.
b)b) Check for happinessCheck for happiness

73. ExamplesExamples
• HFHF
• NN22
• NHNH33
• CHCH44
• CFCF44
• NONO++

74. Exceptions to the OctetExceptions to the Octet
RuleRule

75. Exceptions to the OctetExceptions to the Octet
RuleRule
• Incomplete: An odd number of electrons areIncomplete: An odd number of electrons are
available for bonding. One lone electron is leftavailable for bonding. One lone electron is left
unpaired.unpaired.
• Suboctet: Less than 4 pairs of electrons areSuboctet: Less than 4 pairs of electrons are
assigned to the central atomassigned to the central atom
– Suboctets tend to form coordinateSuboctets tend to form coordinate
covalentbondscovalentbonds
– BHBH33 + NH+ NH33

76. Exceptions to the OctetExceptions to the Octet
RuleRule
• Extended: The central atom has more than 4 pairsExtended: The central atom has more than 4 pairs
of electrons.of electrons.
– At the third energy level and higher, atomsAt the third energy level and higher, atoms
may have empty d orbitals that can be used formay have empty d orbitals that can be used for
bonding.bonding.

77. General RulesGeneral Rules
• The second row elements C, N, O, and F alwaysThe second row elements C, N, O, and F always
obey the octet ruleobey the octet rule
• The second row elements B and Be often haveThe second row elements B and Be often have
fewer than eight electrons around them in theirfewer than eight electrons around them in their
compounds. They are electron deficient and verycompounds. They are electron deficient and very
reactive.reactive.
• The second row elements never exceed the octetThe second row elements never exceed the octet
rule, since their valence orbitals can only hold 8.rule, since their valence orbitals can only hold 8.

78. General RulesGeneral Rules
• Third-row and heavier elements often satisfy theThird-row and heavier elements often satisfy the
octet rule but can exceed the octet rule by usingoctet rule but can exceed the octet rule by using
their empty valence d orbitals.their empty valence d orbitals.
• When writing the Lewis structure for a molecule,When writing the Lewis structure for a molecule,
satisfy the octet rule for the atoms first. Ifsatisfy the octet rule for the atoms first. If
electrons remain after the octet rule has beenelectrons remain after the octet rule has been
satisfied, then place them on the elements havingsatisfied, then place them on the elements having
available d orbitalsavailable d orbitals

79. ResonanceResonance

80. ResonanceResonance
• Resonance is when more than on valid LewisResonance is when more than on valid Lewis
structure can be written for a particularstructure can be written for a particular
molecule. The resulting electron structure ofmolecule. The resulting electron structure of
the molecule is given by the average of thesethe molecule is given by the average of these
resonance structures.resonance structures.

81. ResonanceResonance
• The concept of resonance is necessary becauseThe concept of resonance is necessary because
the localized electron model postulates thatthe localized electron model postulates that
electrons are localized between a given pair ofelectrons are localized between a given pair of
atoms. However, nature does not really operateatoms. However, nature does not really operate
this way. Electrons are really delocalized- theythis way. Electrons are really delocalized- they
move around the entire molecule. The valencemove around the entire molecule. The valence
electrons in a resonance structure distributeelectrons in a resonance structure distribute
themselves equally and produce equal bonds.themselves equally and produce equal bonds.

82. Formal ChargeFormal Charge
Some molecules or polyatomic ions can haveSome molecules or polyatomic ions can have
several non-equivalent Lewis structures.several non-equivalent Lewis structures.
• Example: SOExample: SO44
2-2-
Because of this we assign atomic charges to theBecause of this we assign atomic charges to the
molecules in order to find the right structure.molecules in order to find the right structure.

83. Formal ChargeFormal Charge
The formal charge of an atom in a molecule isThe formal charge of an atom in a molecule is
the difference between the number ofthe difference between the number of
valence electrons on the free atom and thevalence electrons on the free atom and the
number of valence electrons assigned to thenumber of valence electrons assigned to the
atom in the moleculeatom in the molecule
Formal charge = (# of valence electrons onFormal charge = (# of valence electrons on
neutral ‘free atom’) – (# of valence electronsneutral ‘free atom’) – (# of valence electrons
assigned to the atom in the molecule)assigned to the atom in the molecule)

84. Formal ChargeFormal Charge
Assumptions on electron assignment:Assumptions on electron assignment:
• Lone pair electrons belong entirely to theLone pair electrons belong entirely to the
atom in question.atom in question.
• Shared electrons are divided equally betweenShared electrons are divided equally between
the two sharing atoms.the two sharing atoms.

85. Formal Charge ExampleFormal Charge Example
• SOSO44
2-2-
: All single bonds: All single bonds
• Formal charge on each O is -1Formal charge on each O is -1
• Formal charge on S is 2Formal charge on S is 2

86. Formal Charge ExampleFormal Charge Example
• SOSO44
2-2-
: two double bonds, two single: two double bonds, two single
• Formal charge on single bonded O is -1Formal charge on single bonded O is -1
• Formal charge on double bonded O is 0Formal charge on double bonded O is 0
• Formal charge on S is 0Formal charge on S is 0

87. Formal ChargesFormal Charges
1.1. Atoms in molecules try to achieve formal chargesAtoms in molecules try to achieve formal charges
as close to zero as possible.as close to zero as possible.
2.2. Any negative formal charges are expected toAny negative formal charges are expected to
reside on the most electronegative atoms.reside on the most electronegative atoms.
If nonequivalent Lewis structures exist for a species,If nonequivalent Lewis structures exist for a species,
those with formal charges closest to zero andthose with formal charges closest to zero and
with any negative formal charges on the mostwith any negative formal charges on the most
electronegative atoms are considered to bestelectronegative atoms are considered to best
describe the bonding in the molecule or ion.describe the bonding in the molecule or ion.

88. Molecular Structure:Molecular Structure:
The VSEPR ModelThe VSEPR Model

89. VSEPRVSEPR
Valence shell electron repulsion model is useful inValence shell electron repulsion model is useful in
predicting the geometries of molecules formedpredicting the geometries of molecules formed
from nonmetals.from nonmetals.
• The structure around a given atom isThe structure around a given atom is
determined principally by minimizing electrondetermined principally by minimizing electron
– pair repulsion.– pair repulsion.

90. VSEPRVSEPR
• From the Lewis structure, count the electronFrom the Lewis structure, count the electron
pairs around the central atom.pairs around the central atom.
• Lone pairs require more room than bonding pairsLone pairs require more room than bonding pairs
and tend to compress the angles between theand tend to compress the angles between the
bonding pairs.bonding pairs.
• Multiple bonds should be counted as oneMultiple bonds should be counted as one
effective pair.effective pair.
• With a molecule with resonance, all structuresWith a molecule with resonance, all structures
should yield the same shape.should yield the same shape.

91. LinearLinear
• 180°180°

92. Trigonal PlanerTrigonal Planer
• 120°120°

93. TetrahedralTetrahedral
• 109.5°109.5°

94. Trigonal PyramidalTrigonal Pyramidal
• 107°107°

95. Bent/VBent/V
• 104.5°104.5°

96. Tetrahedral ArrangementsTetrahedral Arrangements

97. BipyramidalBipyramidal
ArrangementsArrangements

98. OctahedralOctahedral
ArrangementsArrangements

99. Molecules without aMolecules without a
central atomcentral atom
• The molecular structure of more complicatedThe molecular structure of more complicated
atoms can be predicted from theatoms can be predicted from the
arrangement of pairs around the centerarrangement of pairs around the center
atoms. A combination of shapes will resultatoms. A combination of shapes will result
that allows for minimum repulsionthat allows for minimum repulsion
throughout.throughout.

100. Molecules without aMolecules without a
central atomcentral atom

102. The EndThe End