Writing Lewis Formulas: The Octet Rule

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Writing Lewis Formulas The Octet Rule

• The octet rule states that representative
  elements usually attain stable noble gas electron
  configurations in most of their compounds.
• Lewis dot formulas are based on the octet rule.
• We need to distinguish between bonding (or
  shared) electrons and nonbonding (or unshared or
  lone pairs) of electrons.

• 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.

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Writing Lewis Dot Formulas
N ever Have a Full Octet
Always Have a Full Octet
Sometimes Have a Full Octet
Sometimes Exceed a Full Octet
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Writing Lewis Formulas The Octet Rule

• 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.

• Select a reasonable skeleton
• The least electronegative is the central atom
• Carbon makes 2,3, or 4 bonds
• Nitrogen makes 1(rarely), 2,3, or 4 bonds
• Oxygen makes 1, 2(usually), or 3 bonds
• Oxygen bonds to itself only as O2 or O3,
  peroxides, or superoxides
• Ternary acids (those containing 3 elements)
  hydrogen bonds to the oxygen, not the central
  atom, except phosphates
• For ions or molecules with more than one central
  atom the most symmetrical skeleton is used
• Calculate N, S, and A

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• Count the number of electrons brought to the
  party ( of element times group number)
• For ions we must adjust the number of electrons
  available.
• Add one e- to A for each negative charge.
• Subtract one e- from A for each positive charge.
• Select a reasonable skeleton
• The least electronegative is the central atom
• See prior periodic table for number of electrons
  involved in bonding
• Group I 2 electrons or 1 bond
• Group II 4 electrons or up to 2 bonds
• Group III Al and B, 6 or 8 electrons up to 3 or 4
  bonds
• C,N,O,F must have 8 electrons (up to 4 bonds for
  C, 3 for N, 2 for O, and 1 bond for F).
• All others must have at least 8 electrons (up to
  4 bonds), but may have more.
• The central atom in a molecule or polyatomic ion
  is determined by
• For ions or molecules with more than one central
  atom the most symmetrical skeleton is used
• 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.
• Calculate Formal charges, adjust bonds for lowest
  numbers (zero preferred) and allow for resonance
  structures

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Writing Lewis FormulasThe Octet Rule

• Write Lewis dot and dash formulas for hydrogen
  cyanide, HCN.

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Writing Lewis FormulasThe Octet Rule

• Write Lewis dot and dash formulas for the sulfite
  ion, SO32-.

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Writing Lewis FormulasThe Octet Rule

• 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?

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Resonance

• Write Lewis dot and dash formulas for sulfur
  trioxide, SO3.

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Resonance

• There are three possible structures for SO32-.
• The double bond can be placed in one of three
  places.

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

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Writing Lewis FormulasLimitations of the Octet
Rule

• Write dot and dash formulas for BBr3.

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Writing Lewis FormulasLimitations of the Octet
Rule

• Write dot and dash formulas for AsF5.

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Stereochemistry

• Stereochemistry is the study of the three
  dimensional shapes of molecules.

• Valence Shell Electron Pair Repulsion Theory
• Commonly designated as VSEPR
• Principal originator
• R. J. Gillespie in the 1950s
• Valence Bond Theory
• Involves the use of hybridized atomic orbitals
• Principal originator
• L. Pauling in the 1930s 40s

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• The same basic approach will be used in every
  example of molecular structure prediction

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Polar Molecules The Influence of Molecular
Geometry

• Molecular geometry affects molecular polarity.
• Due to the effect of the bond dipoles and how
  they either cancel or reinforce each other.

A B A
A B A
angular molecule polar
linear molecule nonpolar

• Polar Molecules must meet two requirements
• One polar bond or one lone pair of electrons on
  central atom.
• Neither bonds nor lone pairs can be symmetrically
  arranged that their polarities cancel.

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VSEPR Theory

• Regions of high electron density around the
  central atom are arranged as far apart as
  possible to minimize repulsions.
• There are five basic molecular shapes based on
  the number of regions of high electron density
  around the central atom.

• Lone pairs of electrons (unshared pairs) require
  more volume than shared pairs.
• Consequently, there is an ordering of repulsions
  of electrons around central atom.
• Criteria for the ordering of the repulsions

• Lone pair to lone pair is the strongest
  repulsion.
• Lone pair to bonding pair is intermediate
  repulsion.
• Bonding pair to bonding pair is weakest
  repulsion.
• Mnemonic for repulsion strengths
• lp/lp gt lp/bp gt bp/bp
• Lone pair to lone pair repulsion is why bond
  angles in water are less than 109.5o.

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VSEPR Theory

• Frequently, we will describe two geometries for
  each molecule.
• Electronic geometry is determined by the
  locations of regions of high electron density
  around the central atom(s).
• Molecular geometry determined by the arrangement
  of atoms around the central atom(s).
• Electron pairs are not used in the molecular
  geometry determination just the positions of the
  atoms in the molecule are used.

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VSEPR Theory

• Two regions of high electron density around the
  central atom.

• Three regions of high electron density around the
  central atom.

• Four regions of high electron density around the
  central atom.

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VSEPR Theory

• Five regions of high electron density around the
  central atom.

• Six regions of high electron density around the
  central atom.

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VSEPR Theory

• An example of a molecule that has different
  electronic and molecular geometries is water -
  H2O.
• Electronic geometry is tetrahedral.
• Molecular geometry is bent or angular.

• An example of a molecule that has the same
  electronic and molecular geometries is methane -
  CH4.
• Electronic and molecular geometries are
  tetrahedral.

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Valence Bond (VB) Theory
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Molecular Shapes and Bonding

• In the next sections we will use the following
  terminology
• A central atom
• B bonding pairs around central atom
• U lone pairs around central atom
• For example
• AB3U designates that there are 3 bonding pairs
  and 1 lone pair around the central atom.

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Linear Electronic GeometryAB2 Species (No Lone
Pairs of Electrons on A)
1s 2s 2p Be ?? ??
1s sp hybrid 2p ? ?? ? ?
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Trigonal Planar Electronic Geometry AB3 Species
(No Lone Pairs of Electrons on A)
1s 2s 2p B ??????????????
1s sp2 hybrid ??? ?????? ??? ??? ?
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Tetrahedral Electronic Geometry AB4 Species (No
Lone Pairs of Electrons on A)
2s 2p C He ?????? ? ?? .
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Tetrahedral Electronic Geometry AB4 Species
Valence Bond Theory (Hybridization)
2s 2p C He .
four sp3 hybrids Þ .
Tetrahedral Electronic Geometry AB3U Species
2s 2p N He
four sp3 hybrids Þ
Tetrahedral Electronic Geometry AB2U2 Species

• four sp3 hybrids
• Þ

2s 2p O He
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Tetrahedral Electronic Geometry ABU3 Species
(Three Lone Pairs of Electrons on A)

• Valence Bond Theory (Hybridization)

2s 2p F He ?

• four sp3 hybrids
• Þ ?

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Trigonal Bipyramidal Electronic Geometry AB5,
AB4U, AB3U2, and AB2U3
4s 4p 4d As Ar 3d10 ?? ????????
___ ___ ___ ___ ___
ß five sp3 d hybrids 4d ?? ??
?? ?? ?? ___ ___ ___ ___
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Compounds Containing Double Bonds

• Valence Bond Theory (Hybridization)
• C atom has four electrons.
• Three electrons from each C atom are in sp2
  hybrids.
• One electron in each C atom remains in an
  unhybridized p orbital

2s 2p three sp2 hybrids 2p C ??
?????Þ ??????????? ? ?

• An sp2 hybridized C atom has this shape.
• Remember there will be one electron in each of
  the three lobes.

Top view of an sp2 hybrid
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Compounds Containing Double Bonds

• The single 2p orbital is perpendicular to the
  trigonal planar sp2 lobes.
• The fourth electron is in the p orbital.

Side view of sp2 hybrid with p orbital included.
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Compounds Containing Double Bonds

• Two sp2 hybridized C atoms plus p orbitals in
  proper orientation to form CC double bond.

• The portion of the double bond formed from the
  head-on overlap of the sp2 hybrids is designated
  as a s bond.

• The other portion of the double bond, resulting
  from the side-on overlap of the p orbitals, is
  designated as a p bond.

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Compounds Containing Triple Bonds

• A ? bond results from the head-on overlap of two
  sp hybrid orbitals.

The unhybridized p orbitals form two p bonds.

• Note that a triple bond consists of one ? and
  two p bonds.

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Summary of Electronic Molecular Geometries