Bonding: General Concepts

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Chapter 8

• Bonding General Concepts

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Topics

• Types of chemical bonds
• Electronegativity
• Bond polarity and dipole moments
• Ions Electron configuration and sizes
• Energy effects in binary ionic compounds
• Partial ionic Character of covalent bonds
• The covalent chemical bond A model
• Covalent bond energies and chemical reactions
• The localized electron bonding model
• Lewis structures

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8.1 Types of Chemical bonds

• What is a bond?
• A bond is a force that holds atoms together and
  make them function together
• Bond energy the energy required to break a
  bond.
• Why are compounds (atoms aggregate with each
  other) formed?
• Because this situation gives the system the
  lowest possible energy.

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Ionic Bonding

• An atom with a low ionization energy (metal that
  looses its electron easily) reacts with an atom
  with high electron affinity (nonmetal).
• The electron moves from one atom to the other
  every atom becomes an ion with a certain charge
• Opposite charges hold the ions together.
• Closely packed oppositely charged ions are held
  by strong electrostatic attraction forces

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When atoms lose or gain electrons, they acquire a
noble gas configuration, but do not become noble
gases

• Ionic crystals have great thermal stability and
• consequently acquire high melting points

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Ionic Bond

• When metals and nonmetals
• combine, valence electrons usually are
  transferred from the metal to the nonmetal atoms
  giving rise to electrostatic attraction force
  called ionic bond

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Energy of interaction between a pair of ions
Coulomb's Law

• E 2.31 x 10-19 J nm
• Q is the charge.
• r is the distance between the centers.
• If charges are opposite, E is negative
• It is an exothermic process the ion pair has
  less energy than separated ions
• Same charge, positive E, requires energy to bring
  them together.

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What about identical atoms?How do bonds form
between these atoms?

• The electrons in each atom are attracted to the
  nucleus of the other.
• The electrons repel each other,
• The nuclei repel each other.
• They reach a distance with the lowest possible
  energy.
• The distance between the nuclei is the bond
  length.

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How does a bonding force develop between two
identical atoms?
Atoms are infinitely apart
Energy
0
Zero interacting energy
Internuclear Distance
10
Energy
0
Energy as a function of Internuclear distance
Internuclear Distance
11
Energy
0
Internuclear Distance
12
Energy
0
Most stable state
Internuclear Distance
13
Energy
0
Bond Length
Internuclear Distance
14
Energy
Bond Energy
0
The molecule is more stable than the two
separate atoms
Internuclear Distance
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Covalent Bonding

• Covalent bonding electrons are shared equally
  between two identical atoms.
• on the other extreme (ionic bonding) electrons
  transfer completely to form oppositely charged
  ions
• In between are polar covalent bonds.
• The electrons are not shared evenly.
• One end is slightly positive, the other is
  slightly negative.
• Bond polarity is indicated using small delta ?.

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Polar covalent bond
The density of electron cloud is shifted (some
what) towards one of the two bonded atoms
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-

Effect of electric field on HF molecules
18
-

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8.2 Electronegativity

• The ability of an atom in a molecule to attract
  shared electrons to itself.
• To measure the relative electronegativity,
  imagine an H-X molecule
• Pauling model Relative electronegativities of H
  X are determined by comparing the measured
• H-X bond energy with the expected H-X bond
  energy, which is average of H-H and X-X bond
  energies
• D (H-X) actual - (H-X)expected
• D is expected to be 0 when the two atoms are
  identical

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Electronegativity

• D is known for almost every element
• D gives relative electronegativities of all
  elements.
• D tends to increase left to right.
• D decreases as we go down a group.
• Difference in electronegativity between atoms
  tells about the polarity of the bond

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Electronegativity
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Paulings Electronegativities
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Electronegativity difference
Bond Type
Zero
Covalent
Covalent Character decreases Ionic Character
increases
Intermediate
Large
Relationship between electronegativity and bond
type
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Electronegativity difference bond type
Two identical atoms have the same
electronegativity and share a bonding electron
pair equally. This is called a nonpolar covalent
bond
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8.3 Bond polarity and dipole moments

• A molecule with a center of negative charge and a
  center of positive charge is dipolar (two poles),
  (H-F)
• or has a dipole moment.
• Dipoles will line up in the presence of an
  electric field.

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Representation of the dipolar character

• Any diatomic (two atoms) molecule
• with a polar bond will show a molecular
• dipole moment

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Polar Covalent Bond
A covalent bond with greater electron density
around one of the two atoms
electron rich region
electron poor region
e- rich
e- poor
d
d-
d-
d
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Polarity of polyatomic molecules

• The effect of polar bonds on the polarity of the
  entire molecule depends on the molecule shape
• carbon dioxide has two polar bonds, and is linear
  nonpolar molecule

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Polar molecules

• The effect of polar bonds on the polarity of the
  molecule depends on the molecular shape
• water has two polar bonds and a bent shape the
  highly electronegative oxygen pulls the e- away
  from H very polar!

Thus, H2O molecule has a dipole moment
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Molecules with polar bonds but no resulting
dipole moment
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How to decide for molecular polarity?

• Any diatomic molecule with a polar bond is a
  polar molecule
• For a three or more atoms molecule there are two
  considerations
• There must be a polar bond.
• Geometry cant cancel it out.

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Geometry and polarity

• Three shapes will cancel them out.
• Linear

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• Planar triangles

120º
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• Tetrahedral

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• Others dont cancel, e.g.,
• Bent molecule

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8.4 Ions Electron configuration and size

• Atoms tend to react to form noble gas
  configuration.
• Nonmetals gain electrons from metals or share
  electrons with other nonmetals
• Metals lose electrons to form cations
• Nonmetals can share electrons in covalent bonds.
• Or they can gain electrons to form anions.

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Formation of Cations

• Na 1s22s22p63s1 1 valence electron
• Na1 1s22s22p6 This is a noble gas
  configuration with 8 electrons in the outer
  (valence) level.
• In forming ions, atoms tend to achieve noble gas
  configuration (octet rule or rule of 2)
• Each noble gas (except He 2 e-) has 8 electrons
  in the outer level
• For H it is duet rule (Rule of 2)

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Ionic bond and the octet rule
He
Ne
1s22s1
1s22s22p5
1s2
1s22s22p6
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Noble gas configuration (Octet) rule and covalent
bonding
Atoms are bonded with the electron configuration
of a noble gas that is, the atoms obey the octet
rule
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Predicting formulas of Ionic Compounds

• When the term ionic compound is used it means
  solid state (crystalline) of that compound
• Ions align themselves to maximize attractions
  between opposite charges,
• and to minimize repulsion between like ions.

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Predicting formulas of ionic compounds
Stoichiometry is an important consideration
MgO
1s22s22p63s2
1s22s22p4
Li2O
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Predicting formulas

• Predict the formula of the compound formed from
  Al and O
• Al Ne3s23p1 (should lose 3e- Ne)
• O He2s22p4 ) should gain 2e- Ne)
• The compound to be electrically neutral it has to
  be Al2O3

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Common ions with noble gas configurations in
ionic compounds
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Size of ions

• Ion size increases down a group.
• Cations are smaller than the atoms they came
  from.
• Anions are larger than atoms they came from.
• across a row ions get smaller, and then suddenly
  larger.
• First half are cations.
• Second half are anions.

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Periodic Trends

• Across the period, the change is complicated
  because of the change from predominantly metals
  on the left to nonmetals on the right.

N3-
O2-
F-
B3
Li
C4
Be2
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Atomic and Ionic Radii
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Isoelectronic ions

• Iso - same
• Iso electronic ions have the same of electrons
• Al3 , Mg2, Na, Ne, F-, O2- and N3-
• All have 10 electrons.
• All have the same configuration 1s22s22p6

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Isoelectronic Configurations
Elements that all have the same number of
electrons
For isoelectronic species, the greater the
nuclear charge, the smaller the species
Effective nuclear charge
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Size of Isoelectronic ions

• Positive ions have more protons
• so they are smaller.

Atomic numbers
7
8
9
10
N3-
11
13
O2-
F1-
Ne
Na1
Al3
Mg2
12
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8.5 Energy effects in binary ionic compounds

• Ionic solid is formed because the aggregated
  oppositely charged ions have a lower energy than
  the original elements
• How strongly the ions attract each other in the
  solid state is expressed by the lattice energy
• Lattice energy
• The change in energy that takes place when
  separated gaseous ions are packed together to
  form an ionic solid
• M(g) X-(g) MX

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Lattice energy

• It is usually defined as the energy released when
  ionic solid is formed from its gaseous ions (it
  has ve sign)
• M(g) X-(g) MX(s)
• Lattice energy is a quantitative measure of the
  stability of ionic compound
• The higher the lattice energy the more stable the
  compound

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Na(s) ½F2(g) NaF(s)
Energy changes associated with the formation of
ionic solid

• First sublime Na Na(s) Na(g) DH 109
  kJ/mol
• Ionize Na(g) Na(g) Na(g) e- DH 495
  kJ/mol
• Break F-F Bond ½F2(g) F(g) DH 77
  kJ/mol
• Add electron to F F(g) e- F-(g)
  DH -328 kJ/mol
• Formation of NaF from Na(g) F-(g)
• (Lattice energy) Na(g) F-(g) NaF(s)
• DH -1281 kJ/mol

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Na(s) ½F2(g) NaF(s)

• Overall energy change Na(s) ½F2(g)
  NaF(s) DH -928 kJ/mol

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Calculating Lattice Energy

• k is a constant that depends on the structure of
  the crystal electron configuration of ions
• r is internuclear distance (r radius of
  cationradius of anion)
• Lattice energy is greater with more highly
  charged ions and distances between ions decrease

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8.6 Partial Ionic Character of covalent bonds

• There are probably no totally ionic bonds between
  individual atoms.
• Calculate ionic character of a bond
• Compare measured dipole moment of X-Y bonds to
  the calculated dipole moment of XY- the
  completely ionic case in the gaseous phase

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75
Ionic Character
50
25
Electronegativity difference
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The Relationship Between the Ionic Character of a
Covalent Bond and the Electronegativity
Difference of the Bounded Atoms
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What are the ionic compounds?

• If bonds cant be ionic, what are ionic
  compounds?
• What about polyatomic ions?
• An ionic compound will be defined as any
  substance that conducts electricity when melted

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8.7 The covalent chemical bond A model

• Bonds are the forces that cause a group of atoms
  to behave as a unit.
• Why do chemical bonds occur?
• Due to the tendency of atoms in a system to
  achieve its lowest possible energy
• It takes 1652 kJ to dissociate a mole of CH4 into
  its separate atoms C H.
• Thus, CH4 is stable relative to its separated
  atoms

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• To explain stability, chemical bond term was
  introduced
• Since each hydrogen in CH4 is hooked to the
  carbon, we get the average energy 413 kJ/mol
  Bond Energy
• CH3Cl has 3 C-H, and 1 C Cl and a stabilization
  energy of 1578 kJ/mol
• Thus, C-Cl bond can be calculated as 339 kJ/mol
• The bond is a human invention.
• It is a method of explaining the energy change
  associated with forming molecules.
• Bonds dont exist in nature, but are useful.
• We have a model of a bond.

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8.8 Covalent bond energies and chemical reactions

• Bond energies (BE) is ?H when 1 mole of bonds is
  broken in the gaseous state
• H2(g) 2H(g) ?H 436 kJ
• Cl2(g) 2Cl(g) ?H 243 kJ
• BE is always positive (endothermic)
• Energy is given off when bonds are formed
• H(g) F(g) HF(g) ?H
  -565 kJ
• Bond energy increases with bond polarity
• C-F gt N-F gt O-F gt F-F

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Covalent bond energies and chemical reactions

• Consider stepwise decomposition of CH4
• Each C-H bond has a different energy.
• CH4 CH3 H DH 435 kJ/mol
• CH3 CH2 H DH 453 kJ/mol
• CH2 CH H DH 425 kJ/mol
• CH C H DH 339 kJ/mol
• Each bond is sensitive to its environment but in
  an unsymmetrical way
• Average C-H bond energy 1652/4413kJ/mol

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Multiple bonds

• Single bond one pair of electrons is shared.
• Double bond two pair of electrons are shared.
• triple bond three pair of electrons are shared.
• BE is larger for a multiple bond than for a
  single bond between the same two atoms
• More bonds, shorter bond length.

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Average Bond Energies (kj/mol)
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Lengths of Covalent Bonds
Bond Type Bond Length (pm)
C-C 154
C?C 133
C?C 120
C-N 143
C?N 138
C?N 116
Bond Lengths Triple bond lt Double Bond lt Single
Bond
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Bond Energies and Enthalpy changes in reactions
?H ? D (bonds broken) ?D (bonds formed)
Erequired
Ereleased
Enthalpy change for a reaction
D bond energy per mole of bonds D always has
ve sign
?BE (reactants) ?BE (products)
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Use bond energies to calculate the enthalpy
change for H2 (g) F2 (g) 2HF (g)
DH SBE(reactants) SBE(products)
DH 436.4 156.9 2 x 568.2 -543.1 kJ
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8.9 The localized electron bonding model

• Simple model, easily applied to describe covalent
  bonds
• A molecule is composed of atoms that are bound
  together by sharing pairs of electrons using the
  atomic orbitals of the bound atoms.
• Electron pairs are assumed to be localized on a
  particular atom (Lone pairs) or in the space
  between two atoms (Bonding pairs)

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The Localized Electron Model has three parts

• Description of valence electrons arrangements in
  the molecule using Lewis structures
• Prediction of geometry of the molecule using
  VSEPR (valence shell electron- pair repulsion
  model
• Description of the types of orbitals used by the
  atoms to share electrons or hold lone pairs.

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8.10 Lewis Structure

• Shows how the valence electrons are arranged
  around atoms in a molecule
• One dot for each valence electron.
• A stable compound has all its atoms with a noble
  gas configuration.
• Hydrogen follows the duet rule.
• The rest follow the octet rule.
• Bonding pair is the one between the symbols.

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Lewis Dot Symbols
Nitrogen, N, is in Group 5A and therefore has 5
valence electrons.
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Rules for Writing Lewis Structures

1. Draw skeletal structure of compound showing what
   atoms are bonded to each other. Put least
   electronegative element in the center.
2. Count total number of valence e-. Add 1 for each
   negative charge. Subtract 1 for each positive
   charge.
3. Use one pair of electrons to form a bond (a
   single line) between each pair of atoms.
4. Arrange the remaining electrons to satisfy an
   octet for all atoms (duet for H), starting from
   outer atoms.
5. If a central atom does not have an octet, move in
   lone pairs to form double or triple bonds on the
   central atom as needed.

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Examples

• Fluorine has seven valence electrons
• 9F 1s22s22p5
• A second atom also has seven
• By sharing electrons
• both end with full orbitals

F
F
8 Valence electrons
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Covalent bonding

• Fluorine has seven valence electrons
• A second atom also has seven
• By sharing electrons
• both end with full orbitals

F
F
8 Valence electrons
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Write the Lewis structure of nitrogen trifluoride
(NF3).
Step 1 N is less electronegative than F, put N
in center
Step 2 Count valence electrons N - 5 (2s22p3)
and F - 7 (2s22p5)
5 (3 x 7) 26 valence electrons
Step 3 Draw single bonds between N and F atoms.
Step 4 Arrange remaining 20 electrons to
complete octets
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Write the Lewis structure of the carbonate ion
(CO32-).
Step 1 C is less electronegative than O, put C
in center
Step 2 Count valence electrons C - 4 (2s22p2)
and O - 6 (2s22p4) -2 charge 2e-
4 (3 x 6) 2 24 valence electrons
Step 3 Draw single bonds between C and O atoms
and complete octet on C and O
atoms.
Step 4 - Arrange remaining 18 electrons to
complete octets
Step 5 The central C has only 6 electrons. Form
a double bond.
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Write Lewis Structure for HCN

• .

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Maximum number of bonds (or atoms) surrounding
the central atom

• Central atom Max bonds
• H 1 H-
• O 2 -O-
• N 3 -N-
• C 4 -C-
• X-(F, Cl, Br, I) 1 X-

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Exceptions to the Octet Rule
The Incomplete Octet
BeH2
BF3
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The expanded octet
Some compounds have expanded valence shells,
which means that the central atom has more than
eight electrons around it
Ve 5 7(5) 40
Ve 6 6(7) 48
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Expanded octet
An expanded valence shell may also need to
accommodate lone-pair electrons as well as
bonding pairs
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Molecules with expanded octet

• Terminal atoms are most often halogens. In few
  cases one or more O-atoms are at the end of the
  molecule
• The central atom is a nonmetal in the 3rd, 4th or
  5th period of the Periodic Table
• 3rd P S Cl
• 4th As Se Br Kr
• 5th Sb Te I Xe
• All atoms have d-orbitals available for bondig
  (3d, 4d, 5d) where the extra pairs of electrons
  are located
• Elements of Period 2, NEVER form compounds with
  expanded octet.

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8.12 Resonance
A resonance structure is one of two or more Lewis
structures for a single molecule that cannot be
represented accurately by only one Lewis
structure.
O
O
O
What are the resonance structures of the
carbonate (CO32-) ion?
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• In truth, the electrons that form the second CO
  bond in the double bonds do not always sit
  between that C and that O, but rather can move
  among the three oxygen atoms and the carbon.
• They are not localized, but rather are
  delocalized.

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Molecules that dont follow the octet rule
Molecules with an odd number of valence electrons
have at least one of them unpaired and are called
free radicals
Not to be discused here
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Formal Charge

• For molecules and polyatomic ions that exceed the
  octet there are several different structures.
• Use charges on atoms to help decide which one is
  the real molecule.
• Trying to use the oxidation numbers to put
  charges on atoms in molecules

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Formal charge is the difference between the
number of valence electrons on the free atom and
that assigned to the atom in the molecule.

• Cf Nv (Nu ½ Nb)
• Cf formal charge
• Nv valence e- in the un-bonded atom
• Nu unshared e- owned by the atom
• Nb bonding e- shared by the atom
• In molecules Cf is close to zero
• In ions should be equal to the charge on the
  ion

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Calculation of Formal Charge
90
Using the assumption of formal charges to
evaluate Lewis structure

• Atoms in molecules try to achieve as low a formal
  charge (as close to zero) as possible
• Negative formal charges are expected to be found
  on the most electronegative elements.

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Which is the most likely Lewis structure for CH2O?
?
9.7
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8.13 Molecular Structure VSEPR

• Lewis structures tell us how the atoms are
  connected to each other.
• They dont tell us anything about shape.
• The shape of a molecule can greatly affect its
  properties.
• Valence Shell Electron Pair Repulsion Theory
  allows us to predict geometry

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VSEPR

• Molecules take a shape that puts electron pairs
  as far away from each other as possible.
• The electron-pairs surrounding an atom (valence
  electrons) repel one another and are oriented as
  far apart as possible
• Structure around a given atom is determined
  principally by minimizing electron pair
  repulsion

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• To determine electron pairs Lewis structure
  should be drawn
• Find bonding and nonbonding lone pairs
• Lone pair take more space.
• Multiple bonds count as one pair.

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VSEPR

• The number of pairs determines
• bond angles
• primary structure
• The number of atoms determines
• actual shape

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Valence shell electron pair repulsion (VSEPR)
model
Predict the geometry of the molecule from the
electrostatic repulsions between the electron
(bonding and nonbonding) pairs.
AB2
2
0
97
The best arrangement is to place the two pairs
of Be atom on opposite sides arranged
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VSEPR
AB2
2
0
linear
linear
AB3
3
0
99
F
B
F
F
100
VSEPR
AB2
2
0
linear
linear
AB4
4
0
101
H
C
H
H
H
102
VSEPR
AB2
2
0
linear
linear
AB4
4
0
tetrahedral
tetrahedral
AB5
5
0
103
Cl
Cl
P
Cl
Cl
Cl
104
VSEPR
AB2
2
0
linear
linear
AB4
4
0
tetrahedral
tetrahedral
AB6
6
0
105
F
F
S
F
F
F
F
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Effect of lone pairs on Geometry
Molecules with unshared (Lone) pairs of electrons

• Unshared pair of electrons (under the
• influence of one nucleus) spreads out over a
• volume larger than a bonding pair (under the
• influence of two nuclei).
• The electron pair geometry is approximately
• same as that observed when only single bonds
• are involved
• The bond angles are either equal to the ideal
• values or little less
• The molecular geometry is quite different
• when lone pairs are involved.
• Molecular geometry refers only to the
• positions of the bonded atoms

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VSEPR
trigonal planar
trigonal planar
AB3
3
0
AB2E
2
1
SO2
lt 120o
109
VSEPR
AB4
4
0
tetrahedral
tetrahedral
AB3E
3
1
NH3
lt 109.5o 107o
110
VSEPR
AB4
4
0
tetrahedral
tetrahedral
AB2E2
2
2
H2O
lt 109.5o 104.5o
111
VSEPR
trigonal bipyramidal
trigonal bipyramidal
AB5
5
0
AB4E
4
1
SF4
See-saw
90o, 120o, 180o
112
VSEPR
trigonal bipyramidal
trigonal bipyramidal
AB5
5
0
AB3E2
3
2
ClF3
90o, 180o
113
VSEPR
trigonal bipyramidal
trigonal bipyramidal
AB5
5
0
AB2E3
2
3
I3
180o
114
VSEPR
AB5E
5
1
BrF5
90o, 180o
115
VSEPR
AB4E2
4
2
XeF4
90o, 180o
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The VSEPR model and multiple bond

• For the geometry purposes
• A multiple bond behaves exactly as
• if it were a single electron-pair

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Molecules containing no single central atom

• The central atoms of the molecule should be
• labeled first.
• Geometry can be predicted by focusing on each
• central atom by counting the electron pairs
• around each central atom.

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Predicting Molecular Geometry

1. Draw Lewis structure for molecule.
2. Count number of lone pairs on the central atom
   and number of atoms bonded to the central atom.
3. Use VSEPR to predict the geometry of the molecule.

What are the molecular geometries of SO2 and SF4?
AB4E
AB2E
distorted tetrahedron
bent