Bonding and Molecular Structure

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Chapter 9 and 10

• Bonding and Molecular Structure

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Valence and Core Electrons

• Valence Electrons electrons that participate in
  bonding
• Core Electrons electrons in an atom that do not
  participate in bonding
• Main Group Elements s and p orbitals
• Transition Elements ns and (n-1) d orbitals

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Lewis Dot Symbols for Elements

• Element's symbol represents the atomic nucleus
  together with the core electrons.
• Valence electrons are represented by dots that
  are placed around the symbol.
• These symbols emphasize the ns2np2 octet that
  all noble gases except helium possess.

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Lewis Dot Symbols for Main Group Elements
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Chemical Bond Formation Ionic Bonds

• Electrons are strongly displaced toward one atom
  and away from the other.
• Generally involve metals from the left side of
  the periodic table interacting with nonmetals
  form the far right side.

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Lew Symbols and the Octet Rule for Ionic
Compounds

• The electron configuration of many substances
  after ion formation is that of an noble gas ?
  octet rule.
• Octet rule Main-group elements gain, lose, or
  share in chemical bonding so that they attain a
  valence octet (eight electrons in an atoms
  valence shell).

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Example 1

• The electron configuration of each reactant in
  the formation of KCl gives
• K is that of Ar
• Cl? is also that of Ar.
• The other electrons in the atom are not as
  important in determining the reactivity of that
  substance.
• The octet rule is particularly important in
  compounds involving nonmetals.

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Energy in Ionic Bonding

• When potassium and chlorine atoms approach each
  other we have
• K(g)? K(g) e? Ei 418 kJ
• Cl(g) e?? Cl?(g) Eea ?349 kJ
• K(g)Cl(g)? K(g) Cl?(g) ?E 69 kJ

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Driving Force of Ionic

• Positive ?E energy absorbed ? energetically not
  allowed.
• Driving force must be the formation of the
  crystalline solid.
• K(g) Cl?(g) ? KCl(s)

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Born-Haber Cycle and Lattice Energies

• Overall energetics for the formation of
  crystalline solids can be determined from a
  Born-Haber cycle which accounts for all of the
  steps towards the formation of solid salts from
  the elements. For the formation of KCl from its
  elements we have

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Born-Haber Cycle

• Net energy change of ?434 kJ/mol indicates
  energetically favored.
• Energy for the fifth step is the negative of the
• lattice energy energy required to break ionic
  bonds and sublime (always positive).

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Example 2

• Determine the lattice energy of BaCl2 if the
  heat of sublimation of Ba is 150.9 kJ/mol and the
  1st and 2nd ionization energies are 502 and 966
  kJ/mol, respectively. The heat for the synthesis
  of BaCl2(s) from its elements is ?806.06 kJ/mol.

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LATTICE ENERGIES AND PERIODICITY

• Lattice energy can also be determined from
  Coulombs law
• Directly proportional to charge on each ion.
• Inversely proportional to size of compound (sum
  of ionic radii).

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Lattice Energies

• Table (right) presents the lattice energies for
  alkali and alkaline earth ionic compounds. The
  lattice energies
• decrease for compounds of a particular cation
  with atomic number of the anion.
• decrease for compounds of a particular anion with
  atomic number of the cation.

Lat. E, kJ/mol Lat. E, kJ/mol
LiF 1030 MgCl2 2326
LiCl 834 SrCl2 2127
LiI 730 MgO 3795
NaF 910 CaO 3414
NaCl 788 SrO 3217
NaBr 732 ScN 7547
NaI 682
KF 808
KCl 701
KBr 671
CsCl 657
CsI 600
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Covalent Bond Formation

• Electrons involved in the bond are more or less
  evenly distributed between the atoms, and
  electrons are shared by two nuclei.
• Covalent bonding generally occurs between
  nonmetals, elements that lie in the upper right
  corner of the periodic table.

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The Covalent Bond

• Repulsive forces of the electrons offset by the
  attractive forces between the electrons and the
  two nuclei.
• Bonds are characterized in terms of energy and
  bond distance

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Strengths of Covalent Bonds

• Bonds form because their formation produces lower
  energy state than when atoms are separated.
• Breaking bonds increases the overall energy of
  the system. Energy for breaking bonds has a
  positive sign (negative means that energy is
  given off).
• H - H (g) ? 2H(g) DH 436 kJ.

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Ionic vs Covalent Bond

• Ionic compounds have high melting and boiling
  points and tend to be crystalline
• Covalently bound compounds tend to have lower
  melting points since the attractive forces
  between the molecules are relatively weak.

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Lewis Structures

• Lewis structures are used to indicate valence
  electrons and bonds

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Drawing Lewis Structures

• Determine the arrangement of the atoms in the
  compound with respect to each other and draw a
  skeletal structure.
• If the compound is binary, the first element
  written down is usually the central atom
  (hydrogen is an exception to this).
• With a ternary compound (one with three kinds of
  elements) the middles atom in the formula is
  usually the central one.

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Lewis Structures Cont.

• Determine the total number of valence electrons.
• Subtract two electrons from the valence total for
  each bond in the skeletal structure.

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Lewis Structures Cont.

• Determine how many electrons are required for
  each element to have a total of eight (there are
  several exceptions to this rule).
• If a sufficient number of electrons are
  available, distribute the remaining electrons
  around the element symbols.
• If a sufficient number is not available, add
  additional bonds making certain to subtract the
  electrons used. (Multiple bonds often occur
  among the atoms C, N, and O.

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Example 3

• Draw Lewis Electron Dot Structures for the
  following

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Lewis Structures and Resonance

• Quantum theory indicates that any position is
  possible for an electron.
• Equivalent electron positions often possible
• E.g. SO2 OS-O and O-SO.
• Each structure equally likely.
• the true form of the molecule is a hybrid of
  these and is called resonance and the hybrid form
  is called a resonance hybrid.

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Example 4

• Draw Lewis Electron Dot Structures for the
  following

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Exceptions to the Octet Rule

• Although many molecules obey the octet rule,
  there are exceptions where the central atom has
  fewer or more than eight electrons.
• Compounds in which an atom has fewer than eight
  valence electrons

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• Compounds in which an atom has more than eight
  valence electrons
• Generally, if a nonmetal is in the third period
  or greater it can accommodate as many as twelve
  electrons, if it is the central atom.
• These elements have unfilled d subshells that
  can be used for bonding.

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Free Radicals

• Molecules with an odd number of electrons

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Example 5

• Draw Lewis Electron Dot Structures for the
  following XeF4, ICl3, and SF4

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FORMAL CHARGES

• Formal Charge (of an atom in a Lewis formula) the
  hypothetical charge obtained by assuming that
  bonding electrons are equally shared between the
  two atoms involved in the bond. Lone pair
  electrons belong only to the atom to which they
  are bound.
• Formal Charge group number number of lone
  pair electrons ½ bonding electrons

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Example 6

• Determine the formal charge on all elements
  PCl3, PCl5, and HNO3.

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• Formal charge (FC) allows the prediction of the
  more likely resonance structure.
• To determine the more likely resonance structure
• FC should be as close to zero as possible.
• Negative charge should reside on the most
  electronegative and positive charge on the least
  electronegative element.

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Example 7

• Draw the resonance structures of H2SO4
  determine the formal charge on each element and
  decide which is the most likely structure.

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Bond Properties

• Bonds Order
• Bond Length
• Bond Energy
• Bond Polarity

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Bond Order

• The order of a bond is the number of bonding
  electrons pairs shared by two atoms in a
  molecule.
• A fractional bond order is possible in molecules
  and ions having resonance structures

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Bond Length

• Bond length is the distance between the nuclei of
  two bonded atoms. Bond length is determined in a
  large part by the size of the atoms.
• Bond length becomes shorter as bond order
  increases

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Bond Dissociation Enthalpies

• Bond dissociation energy, BE the energy
  required to break one mole of a type of bond in
  an isolated molecule in the gas phase.
• Useful for estimation of heat of unknown
  reactions.

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• Bond energies are all positive
• The energies are average bond energies
• Bond energies are defined in terms of gaseous
  atoms of molecular fragments
• Bond energies increase with bond order

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• Hesss law can be used with bond dissociation
  energies to estimate the enthalpy change of a
  reaction.
• The breaking in a C H bond would be C H(g) ?
  C(g) H(g) ?H BE 410 kJ.
• Sign always positive since energy must be
  supplied to break bond.

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Using Bond Dissociation Enthalpies

• Estimate the heat of formation of H2O(g) from
  bond dissociation energies. Thus determine
• H2(g) ½ O2(g) ? H2O(g) ?
•   H H (g) ? 2H(g) ?H BE 436 kJ 
• ½ OO ? O(g) ?H BE 494/2 247 kJ
• 2H(g) O(g) ? H O H (g) ?H ?2BE ?2459
  kJ  
• H2(g) ½ O2(g) ? H2O(g) ?235 kJ
• Actual ?241.8 KJ
• Can be determined by summing all the energies for
  the bonds broken and subtract from it the sum of
  the energies for the bonds formed.

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Example 8

• Estimate the energy change for the chlorination
  of ethylene
• CH2CH2(g) Cl2(g)? CH2ClCH2Cl

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Bond Polarity and Electronegativity

• Electronegativities
• increase from bottom to top of periodic table and
  
• increase to a maximum towards the top right.
• can provide an insight as to the type of bond
  that would be expected.

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Electronegativities
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Bond Polarity

• Ionic bonds formed when displacement of electrons
  is essentially complete
• Covalent bonds forms when no displacement of
  electrons occurs
• Polar covalent forms when bond pair is not
  equally shared between two atoms and the
  electrons are displace toward one of the atoms
  from a point midway between them

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Polar Bonds

• With a polar bond exists between two atoms, a
  small charge on the atom due to that bond
  develops. ? and ?? designates which is the
  positive and negative side respectively

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Example 9

• Determine the relative polarities of HF, HCl, HBr
  and HI.

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Molecular Geometry and Directional Bonding

• Atoms oriented in very well defined relative
  positions in the molecule.
• Molecular Geometry general shape of the
  molecule as determined by the relative positions
  of the atomic nuclei.

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• Theories describing the structure and bonding of
  molecules are
• VSEPR considers mostly electrostatics in
  determining the geometry of the molecule.
• Valence Bond Theory considers quantum mechanics
  and hybridization of atomic orbitals.
• Molecular Orbital Theory claims that upon bond
  formation new orbitals that are linear
  combinations of the atomic orbitals are formed.

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Valence Share Electron Pair Repulsion (VSEPR)
Theory

• Valence Share Electron-Pair Repulsion (VSEPR)
  model allows us to predict the molecular shape by
  assuming that the repulsive forces of electron
  pairs cause them to be as far apart as possible
  from each other.

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PREDICTING EXPECTED GEOMETRY ACCORDING TO VSEPR
THEORY.

• Lewis dot structure determines the total of
  electrons around the central atom. Multiple
  bonds (double and triple) count as one.
• The number of bonding and nonbonding electron
  pairs determines the geometry of electron pairs
  and the molecular geometry.
• Only the valence electron pairs are considered in
  determining the geometry.

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Procedure

• Count the number of valence electrons and write
  the Lewis Structure
• Determine the number of unshared electron pairs,
  E, and the number of shared pairs, X.
• Determine the arrangement of bonds and unshared
  pairs that minimizes electron pair repulsions
• Describe the shape of the molecule in terms of
  the positions of the atoms
• Note Multiple bonds are treated as single bonds
• Lone e? Pairs affect geometry more than bonding

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Example 10

• BeCl2.and CO2

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Example 11

• BF3, COCl2, O3, SO2

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Example 12

• CH4, PCl3, H2O

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Example 13

• PCl5, SF4, ClF3, XeF2 (lone pair in axial
  position for a trigonal bipyramidal structure).

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Example 14

• SF6, IF5, XeF4

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Polarity of Molecules

• Bond dipole a positive charge next to a negative
  charge.
• Dipole moment, ? the magnitude of the net bond
  dipole of a molecule ? Qxr Q the net charge
  separation r the separation distance. Units
  debyes (D) where 1 D 33.36x10?30 Cm.

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The Dipole     A dipole arises when two
electrical charges of equal magnitude but
opposite sign are separated by distance.
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• A polar bond forms when two atoms of between two
  atoms involved in a bond have significantly
  different electronegativities.
• Most electronegative substance will have a slight
  negative charge (represented as ??)
• The positive (electron poor) side of the bond is
  represented as ? or
• ? points in direction of the negative charge.
• Net polarity (dipole moment) of a molecule is
  obtained using the vector sum of polarities of
  individual bonds.

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For a polyatomic molecule we treat the dipoles as
3D vectors
The sum of these vectors will give us the dipole
for the molecule
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Example 14

• Determine if NH3, H2O, CO2 have dipole moments.

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MOLECULAR SHAPESVALENCE BOND THEORY (VBT)

• Valence Bond Theory a quantum mechanical
  description of bonding that pictures covalent
  bond formation as the overlap of two singly
  occupied atomic orbitals.

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• VSEPR effective but ignores the orbital concepts
  discussed in quantum mechanics.
• H2 forms due to overlap of two 1s orbitals.
• Electron densities from p-subshell electrons
  overlap to produce a bond in F2.

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Overlap of Orbitals
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The degree of overlap is determined by the
systems potential energy
equilibrium bond distance
The point at which the potential energy is a
minimum is called the equilibrium bond distance
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Hybrid Orbitals

• In molecules, the orbitals occupied by electron
  pairs are seldom pure s or p orbitals.
  Instead, they are hybrid orbitals formed by
  combining s and p or s,p and d orbitals.
• Hybridization determined by using VSEPR to
  establish the geometry, i.e., the number of
  electron clouds around the central atom. The
  number of electron clouds the number of hybrid
  orbitals.

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sp
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Formation of sp hybrid orbitals The combination
of an s orbital and a p orbital produces 2 new
orbitals called sp orbitals.
2s
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sp2
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Formation of sp2 hybrid orbitals
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sp3
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Formation of sp3 hybrid orbitals
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sp3d
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sp3d2
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Electron Geometry and Hybrid Orbitals
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Valence Bond Theory and Multiple Bonds   Sigma
bond (s) ? A bond where the line of electron
density is concentrated symmetrically along the
line connecting the two atoms. Only 1 sigma bond
can exist between two atoms
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Pi bond (p) ? A bond where the overlapping
regions exist above and below the internuclear
axis (with a nodal plane along the internuclear
axis). Occur when multiple onds form between two
atoms
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• C2H4 planar with a trigonal geometry sp2
  hybridization for each of the carbon atoms and
  they form ? bonds with hydrogen.
• Each carbon has 4 orbitals in its valence shell.
  This means one of the p-orbitals for each C is
  not hybridized.
• Proximity to each other results in overlap to
  give a charge distribution resembling a cloud
  which is above and below the plane of the
  molecule - a ? bond

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H2CCH2
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H2CCH2
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C2H2

• sp (linear) hybridized. Leads to the existence of
  a ? bond as well as two ? bonds.

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Example HC?CH
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Summarizing

• single bond is a ? bond,
• double bond is a ? bond and a ? bond,
• triple bond is a ? bond and 2 ? bonds.

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MO Theory of Bonding

• Molecular Orbital Theory extends quantum theory
  and states that electrons spread throughout the
  molecule in molecular orbitals region in a
  molecule in which an electron is likely to be
  which is similar to the concept discussed in
  quantum theory. Molecular orbitals are
  considered to be the result of the combination of
  atomic orbitals.

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Molecular Orbital (MO) Theory
ANTBONDING
These two new orbitals have different energies. 
BONDING
The one that is lower in energy is called the
bonding orbital, The one higher in energy is
called an antibonding orbital.
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Energy level diagrams / molecular orbital diagrams
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MO Theory for 2nd row diatomic molecules   Molecul
ar Orbitals (MOs) from Atomic Orbitals
(AOs)   1. of Molecular Orbitals of Atomic
Orbitals 2. The number of electrons occupying
the Molecular orbitals is equal to the sum of the
valence electrons on the constituent atoms. 3.
When filling MOs the Pauli Exclusion Principle
Applies (2 electrons per Molecular
Orbital) 4. For degenerate MOs, Hund's rule
applies. 5. AOs of similar energy combine more
readily than ones of different energy 6. The
more overlap between AOs the lower the energy of
the bonding orbital they create and the higher
the energy of the antibonding orbital.
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Example Li2