Chapter%209%20Chemical%20Bonding%20I:%20Lewis%20Theory

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Chapter 9Chemical Bonding ILewis Theory
Chemistry A Molecular Approach, 1st Ed.Nivaldo
Tro
Roy Kennedy Massachusetts Bay Community
College Wellesley Hills, MA
2008, Prentice Hall
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Bonding Theories

• explain how and why atoms attach together
• explain why some combinations of atoms are stable
  and others are not
• why is water H2O, not HO or H3O
• one of the simplest bonding theories was
  developed by G.N. Lewis and is called Lewis
  Theory
• Lewis Theory emphasizes valence electrons to
  explain bonding
• using Lewis Theory, we can draw models called
  Lewis structures that allow us to predict many
  properties of molecules
• aka Electron Dot Structures
• such as molecular shape, size, polarity

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Why Do Atoms Bond?

• processes are spontaneous if they result in a
  system with lower potential energy
• chemical bonds form because they lower the
  potential energy between the charged particles
  that compose atoms
• the potential energy between charged particles is
  directly proportional to the product of the
  charges
• the potential energy between charged particles is
  inversely proportional to the distance between
  the charges

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Potential Energy Between Charged Particles

• ?0 is a constant
• 8.85 x 10-12 C2/Jm
• for charges with the same sign, Epotential is
  and the magnitude gets less positive as the
  particles get farther apart
• for charges with the opposite signs, Epotential
  is ? and the magnitude gets more negative as the
  particles get closer together
• remember the more negative the potential energy,
  the more stable the system becomes

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Potential Energy BetweenCharged Particles
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Bonding

• a chemical bond forms when the potential energy
  of the bonded atoms is less than the potential
  energy of the separate atoms
• have to consider following interactions
• nucleus-to-nucleus repulsion
• electron-to-electron repulsion
• nucleus-to-electron attraction

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Types of Bonds
Types of Atoms Type of Bond Bond Characteristic
metals to nonmetals Ionic electrons transferred
nonmetals to nonmetals Covalent electrons shared
metal to metal Metallic electrons pooled
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Types of Bonding
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Ionic Bonds

• when metals bond to nonmetals, some electrons
  from the metal atoms are transferred to the
  nonmetal atoms
• metals have low ionization energy, relatively
  easy to remove an electron from
• nonmetals have high electron affinities,
  relatively good to add electrons to

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Covalent Bonds

• nonmetals have relatively high ionization
  energies, so it is difficult to remove electrons
  from them
• when nonmetals bond together, it is better in
  terms of potential energy for the atoms to share
  valence electrons
• potential energy lowest when the electrons are
  between the nuclei
• shared electrons hold the atoms together by
  attracting nuclei of both atoms

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Determining the Number of Valence Electrons in an
Atom

• the column number on the Periodic Table will tell
  you how many valence electrons a main group atom
  has
• Transition Elements all have 2 valence electrons
  Why?

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Lewis Symbols of Atoms

• aka electron dot symbols
• use symbol of element to represent nucleus and
  inner electrons
• use dots around the symbol to represent valence
  electrons
• pair first two electrons for the s orbital
• put one electron on each open side for p
  electrons
• then pair rest of the p electrons

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Lewis Symbols of Ions

• Cations have Lewis symbols without valence
  electrons
• Lost in the cation formation
• Anions have Lewis symbols with 8 valence
  electrons
• Electrons gained in the formation of the anion

Li Li1
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What We Know

• the noble gases are the least reactive group of
  elements
• the alkali metals are the most reactive metals
  and their atoms almost always lose 1 electron
  when they react
• the halogens are the most reactive group of
  nonmetals and in a lot of reactions they gain 1
  electron

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Stable Electron ArrangementsAnd Ion Charge

• Metals form cations by losing enough electrons to
  get the same electron configuration as the
  previous noble gas
• Nonmetals form anions by gaining enough electrons
  to get the same electron configuration as the
  next noble gas
• The noble gas electron configuration must be very
  stable

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Octet Rule

• when atoms bond, they tend to gain, lose, or
  share electrons to result in 8 valence electrons
• ns2np6
• noble gas configuration
• many exceptions
• H, Li, Be, B attain an electron configuration
  like He
• He 2 valence electrons
• Li loses its one valence electron
• H shares or gains one electron
• though it commonly loses its one electron to
  become H
• Be loses 2 electrons to become Be2
• though it commonly shares its two electrons in
  covalent bonds, resulting in 4 valence electrons
• B loses 3 electrons to become B3
• though it commonly shares its three electrons in
  covalent bonds, resulting in 6 valence electrons
• expanded octets for elements in Period 3 or below
  
• using empty valence d orbitals

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Lewis Theory

• the basis of Lewis Theory is that there are
  certain electron arrangements in the atom that
  are more stable
• octet rule
• bonding occurs so atoms attain a more stable
  electron configuration
• more stable lower potential energy
• no attempt to quantify the energy as the
  calculation is extremely complex

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Properties of Ionic Compounds

• hard and brittle crystalline solids
• all are solids at room temperature
• melting points generally gt 300?C
• the liquid state conducts electricity
• the solid state does not conduct electricity
• many are soluble in water
• the solution conducts electricity well

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Conductivity of NaCl
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Lewis Theory and Ionic Bonding

• Lewis symbols can be used to represent the
  transfer of electrons from metal atom to nonmetal
  atom, resulting in ions that are attracted to
  each other and therefore bond

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Predicting Ionic FormulasUsing Lewis Symbols

• electrons are transferred until the metal loses
  all its valence electrons and the nonmetal has an
  octet
• numbers of atoms are adjusted so the electron
  transfer comes out even

Li2O
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Energetics of Ionic Bond Formation

• the ionization energy of the metal is endothermic
• Na(s) ? Na(g) 1 e - DH 603 kJ/mol
• the electron affinity of the nonmetal is
  exothermic
• ½Cl2(g) 1 e - ? Cl-(g) DH - 227 kJ/mol
• generally, the ionization energy of the metal is
  larger than the electron affinity of the
  nonmetal, therefore the formation of the ionic
  compound should be endothermic
• but the heat of formation of most ionic compounds
  is exothermic and generally large Why?
• Na(s) ½Cl2(g) ? NaCl(s) DHf -410 kJ/mol

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Ionic Bonds

• electrostatic attraction is nondirectional!!
• no direct anion-cation pair
• no ionic molecule
• chemical formula is an empirical formula, simply
  giving the ratio of ions based on charge balance
• ions arranged in a pattern called a crystal
  lattice
• every cation surrounded by anions and every
  anion surrounded by cations
• maximizes attractions between and - ions

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Lattice Energy

• the lattice energy is the energy released when
  the solid crystal forms from separate ions in the
  gas state
• always exothermic
• hard to measure directly, but can be calculated
  from knowledge of other processes
• lattice energy depends directly on size of
  charges and inversely on distance between ions

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Born-Haber Cycle

• method for determining the lattice energy of an
  ionic substance by using other reactions
• use Hesss Law to add up heats of other processes
• DHf(salt) DHf(metal atoms, g)
  DHf(nonmetal atoms, g) DHf(cations, g)
  DHf(anions, g) DHf(crystal lattice)
• DHf(crystal lattice) Lattice Energy
• metal atoms (g) ? cations (g), DHf ionization
  energy
• dont forget to add together all the ionization
  energies to get to the desired cation
• M2 1st IE 2nd IE
• nonmetal atoms (g) ? anions (g), DHf electron
  affinity

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Born-Haber Cycle for NaCl
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Practice - Given the Information Below, Determine
the Lattice Energy of MgCl2
Mg(s) Mg(g) DH1f 147.1 kJ/mol ½ Cl2(g)
Cl(g) DH2f 121.3 kJ/mol Mg(g) Mg1(g)
DH3f 738 kJ/mol Mg1(g) Mg2(g) DH4f
1450 kJ/mol Cl(g) Cl-1(g) DH5f -349
kJ/mol Mg(s) Cl2(g) MgCl2(s) DH6f -641.3
kJ/mol
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Practice - Given the Information Below, Determine
the Lattice Energy of MgCl2
Mg(s) Mg(g) DH1f 147.1 kJ/mol ½ Cl2(g)
Cl(g) DH2f 121.3 kJ/mol Mg(g) Mg1(g)
DH3f 738 kJ/mol Mg1(g) Mg2(g) DH4f
1450 kJ/mol Cl(g) Cl-1(g) DH5f -349
kJ/mol Mg(s) Cl2(g) MgCl2(s) DH6f -641.3
kJ/mol
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Trends in Lattice EnergyIon Size

• the force of attraction between charged particles
  is inversely proportional to the distance between
  them
• larger ions mean the center of positive charge
  (nucleus of the cation) is farther away from
  negative charge (electrons of the anion)
• larger ion weaker attraction smaller lattice
  energy

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Lattice Energy vs. Ion Size
Metal Chloride Lattice Energy (kJ/mol)
LiCl -834
NaCl -787
KCl -701
CsCl -657
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Trends in Lattice EnergyIon Charge

• the force of attraction between oppositely
  charged particles is directly proportional to the
  product of the charges
• larger charge means the ions are more strongly
  attracted
• larger charge stronger attraction larger
  lattice energy
• of the two factors, ion charge generally more
  important

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Example 9.2 Order the following ionic compounds
in order of increasing magnitude of lattice
energy.CaO, KBr, KCl, SrO
First examine the ion charges and order by
product of the charges
Ca2 O2-, K Br-, K Cl-, Sr2 O2-
(KBr, KCl) lt (CaO, SrO)
Then examine the ion sizes of each group and
order by radius larger lt smaller
(KBr, KCl) same cation, Br- gt Cl- (same Group)
(CaO, SrO) same anion, Sr2 gt Ca2 (same Group)
KBr lt KCl lt (CaO, SrO)
KBr lt KCl lt SrO lt CaO
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Ionic BondingModel vs. Reality

• ionic compounds have high melting points and
  boiling points
• MP generally gt 300C
• all ionic compounds are solids at room
  temperature
• because the attractions between ions are strong,
  breaking down the crystal requires a lot of
  energy
• the stronger the attraction (larger the lattice
  energy), the higher the melting point

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Ionic BondingModel vs. Reality

• ionic solids are brittle and hard
• the position of the ion in the crystal is
  critical to establishing maximum attractive
  forces displacing the ions from their positions
  results in like charges close to each other and
  the repulsive forces take over

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Ionic BondingModel vs. Reality

• ionic compounds conduct electricity in the liquid
  state or when dissolved in water, but not in the
  solid state
• to conduct electricity, a material must have
  charged particles that are able to flow through
  the material
• in the ionic solid, the charged particles are
  locked in position and cannot move around to
  conduct
• in the liquid state, or when dissolved in water,
  the ions have the ability to move through the
  structure and therefore conduct electricity

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Covalent BondingBonding and Lone Pair Electrons

• Covalent bonding results when atoms share pairs
  of electrons to achieve an octet
• Electrons that are shared by atoms are called
  bonding pairs
• Electrons that are not shared by atoms but belong
  to a particular atom are called lone pairs
• aka nonbonding pairs

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Single Covalent Bonds

• two atoms share a pair of electrons
• 2 electrons
• one atom may have more than one single bond

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Double Covalent Bond

• two atoms sharing two pairs of electrons
• 4 electrons

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Triple Covalent Bond

• two atoms sharing 3 pairs of electrons
• 6 electrons

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Covalent BondingPredictions from Lewis Theory

• Lewis theory allows us to predict the formulas of
  molecules
• Lewis theory predicts that some combinations
  should be stable, while others should not
• because the stable combinations result in
  octets
• Lewis theory predicts in covalent bonding that
  the attractions between atoms are directional
• the shared electrons are most stable between the
  bonding atoms
• resulting in molecules rather than an array

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Covalent BondingModel vs. Reality

• molecular compounds have low melting points and
  boiling points
• MP generally lt 300C
• molecular compounds are found in all 3 states at
  room temperature
• melting and boiling involve breaking the
  attractions between the molecules, but not the
  bonds between the atoms
• the covalent bonds are strong
• the attractions between the molecules are
  generally weak
• the polarity of the covalent bonds influences the
  strength of the intermolecular attractions

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Intermolecular Attractions vs. Bonding
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Ionic BondingModel vs. Reality

• some molecular solids are brittle and hard, but
  many are soft and waxy
• the kind and strength of the intermolecular
  attractions varies based on many factors
• the covalent bonds are not broken, however, the
  polarity of the bonds has influence on these
  attractive forces

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Ionic BondingModel vs. Reality

• molecular compounds do not conduct electricity in
  the liquid state
• molecular acids conduct electricity when
  dissolved in water, but not in the solid state
• in molecular solids, there are no charged
  particles around to allow the material to conduct
• when dissolved in water, molecular acids are
  ionized, and have the ability to move through the
  structure and therefore conduct electricity

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Bond Polarity

• covalent bonding between unlike atoms results in
  unequal sharing of the electrons
• one atom pulls the electrons in the bond closer
  to its side
• one end of the bond has larger electron density
  than the other
• the result is a polar covalent bond
• bond polarity
• the end with the larger electron density gets a
  partial negative charge
• the end that is electron deficient gets a partial
  positive charge

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HF
EN 2.1
EN 4.0
d
d-
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Electronegativity

• measure of the pull an atom has on bonding
  electrons
• increases across period (left to right) and
• decreases down group (top to bottom)
• fluorine is the most electronegative element
• francium is the least electronegative element
• the larger the difference in electronegativity,
  the more polar the bond
• negative end toward more electronegative atom

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Electronegativity Scale
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Electronegativity and Bond Polarity

• If difference in electronegativity between bonded
  atoms is 0, the bond is pure covalent
• equal sharing
• If difference in electronegativity between bonded
  atoms is 0.1 to 0.4, the bond is nonpolar
  covalent
• If difference in electronegativity between bonded
  atoms 0.5 to 1.9, the bond is polar covalent
• If difference in electronegativity between bonded
  atoms larger than or equal to 2.0, the bond is
  ionic

100
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Bond Polarity
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Bond Dipole Moments

• the dipole moment is a quantitative way of
  describing the polarity of a bond
• a dipole is a material with positively and
  negatively charged ends
• measured
• dipole moment, m, is a measure of bond polarity
• it is directly proportional to the size of the
  partial charges and directly proportional to the
  distance between them
• m (q)(r)
• not Coulombs Law
• measured in Debyes, D
• the percent ionic character is the percentage of
  a bonds measured dipole moment to what it would
  be if full ions

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Dipole Moments
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Water a Polar Molecule
stream of water attracted to a charged glass rod
stream of hexane not attracted to a charged glass
rod
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Example 9.3(c) - Determine whether an N-O bond is
ionic, covalent, or polar covalent.

• Determine the electronegativity of each element
• N 3.0 O 3.5
• Subtract the electronegativities, large minus
  small
• (3.5) - (3.0) 0.5
• If the difference is 2.0 or larger, then the bond
  is ionic otherwise its covalent
• difference (0.5) is less than 2.0, therefore
  covalent
• If the difference is 0.5 to 1.9, then the bond is
  polar covalent otherwise its covalent
• difference (0.5) is 0.5 to 1.9, therefore polar
  covalent

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Lewis Structures of Molecules

• shows pattern of valence electron distribution in
  the molecule
• useful for understanding the bonding in many
  compounds
• allows us to predict shapes of molecules
• allows us to predict properties of molecules and
  how they will interact together

57
Lewis Structures

• use common bonding patterns
• C 4 bonds 0 lone pairs, N 3 bonds 1 lone
  pair, O 2 bonds 2 lone pairs, H and halogen
  1 bond, Be 2 bonds 0 lone pairs, B 3 bonds
  0 lone pairs
• often Lewis structures with line bonds have the
  lone pairs left off
• their presence is assumed from common bonding
  patterns
• structures which result in bonding patterns
  different from common have formal charges

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Writing Lewis Structures of Molecules HNO3

• Write skeletal structure
• H always terminal
• in oxyacid, H outside attached to Os
• make least electronegative atom central
• N is central

• Count valence electrons
• sum the valence electrons for each atom
• add 1 electron for each - charge
• subtract 1 electron for each charge

N 5 H 1 O3 36 18 Total 24 e-
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Writing Lewis Structures of Molecules HNO3

1. Attach central atom to the surrounding atoms with
   pairs of electrons and subtract from the total

Electrons Start 24 Used 8 Left 16
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Writing Lewis Structures of Molecules HNO3

• Complete octets, outside-in
• H is already complete with 2
• 1 bond

and re-count electrons
N 5 H 1 O3 36 18 Total 24 e-
Electrons Start 24 Used 8 Left 16
Electrons Start 16 Used 16 Left 0
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Writing Lewis Structures of Molecules HNO3

• If all octets complete, give extra electrons to
  central atom.
• elements with d orbitals can have more than 8
  electrons
• Period 3 and below
• If central atom does not have octet, bring in
  electrons from outside atoms to share
• follow common bonding patterns if possible

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Practice - Lewis Structures

• CO2
• SeOF2
• NO2-1

• H3PO4
• SO3-2
• P2H4

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Practice - Lewis Structures

• CO2
• SeOF2
• NO2-1

• H3PO4
• SO3-2
• P2H4

16 e-
32 e-
26 e-
26 e-
18 e-
14 e-
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Formal Charge

• during bonding, atoms may wind up with more or
  less electrons in order to fulfill octets - this
  results in atoms having a formal charge
• FC valence e- - nonbonding e- - ½ bonding e-
• left O FC 6 - 4 - ½ (4) 0
• S FC 6 - 2 - ½ (6) 1
• right O FC 6 - 6 - ½ (2) -1
• sum of all the formal charges in a molecule 0
• in an ion, total equals the charge

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Writing Lewis Formulas of Molecules (contd)

• Assign formal charges to the atoms
• formal charge valence e- - lone pair e- - ½
  bonding e-
• follow the common bonding patterns

0
1
-1
all 0
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Common Bonding Patterns
-
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Practice - Assign Formal Charges

• CO2
• SeOF2
• NO2-1

• H3PO4
• SO3-2
• P2H4

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Practice - Assign Formal Charges
-1

• CO2
• SeOF2
• NO2-1

• H3PO4
• SO3-2
• P2H4

P 1 rest 0
all 0
-1
-1
Se 1
-1
S 1
-1
-1
all 0
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Resonance

• when there is more than one Lewis structure for a
  molecule that differ only in the position of the
  electrons, they are called resonance structures
• the actual molecule is a combination of the
  resonance forms a resonance hybrid
• it does not resonate between the two forms,
  though we often draw it that way
• look for multiple bonds or lone pairs

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Resonance
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Ozone Layer
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Rules of Resonance Structures

• Resonance structures must have the same
  connectivity
• only electron positions can change
• Resonance structures must have the same number of
  electrons
• Second row elements have a maximum of 8 electrons
• bonding and nonbonding
• third row can have expanded octet
• Formal charges must total same
• Better structures have fewer formal charges
• Better structures have smaller formal charges
• Better structures have - formal charge on more
  electronegative atom

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Drawing Resonance Structures

1. draw first Lewis structure that maximizes octets
2. assign formal charges
3. move electron pairs from atoms with (-) formal
   charge toward atoms with () formal charge
4. if () fc atom 2nd row, only move in electrons if
   you can move out electron pairs from multiple
   bond
5. if () fc atom 3rd row or below, keep bringing in
   electron pairs to reduce the formal charge, even
   if get expanded octet.

-1
-1
1
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Exceptions to the Octet Rule

• expanded octets
• elements with empty d orbitals can have more than
  8 electrons
• odd number electron species e.g., NO
• will have 1 unpaired electron
• free-radical
• very reactive
• incomplete octets
• B, Al

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Drawing Resonance Structures

1. draw first Lewis structure that maximizes octets
2. assign formal charges
3. move electron pairs from atoms with (-) formal
   charge toward atoms with () formal charge
4. if () fc atom 2nd row, only move in electrons if
   you can move out electron pairs from multiple
   bond
5. if () fc atom 3rd row or below, keep bringing in
   electron pairs to reduce the formal charge, even
   if get expanded octet.

-1
2
-1
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Practice - Identify Structures with Better or
Equal Resonance Forms and Draw Them
-1

• CO2
• SeOF2
• NO2-1

• H3PO4
• SO3-2
• P2H4

P 1
all 0
-1
-1
Se 1
-1
S 1
-1
-1
all 0
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Practice - Identify Structures with Better or
Equal Resonance Forms and Draw Them

• CO2
• SeOF2
• NO2-1

• H3PO4
• SO3-2
• P2H4

-1
all 0
none
1
-1
all 0
S 0 in all res. forms
1
-1
-1
none
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Bond Energies

• chemical reactions involve breaking bonds in
  reactant molecules and making new bond to create
  the products
• the DHreaction can be calculated by comparing
  the cost of breaking old bonds to the profit from
  making new bonds
• the amount of energy it takes to break one mole
  of a bond in a compound is called the bond energy
• in the gas state
• homolytically each atom gets ½ bonding electrons

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Trends in Bond Energies

• the more electrons two atoms share, the stronger
  the covalent bond
• CC (837 kJ) gt CC (611 kJ) gt C-C (347 kJ)
• CN (891 kJ) gt CN (615 kJ) gt C-N (305 kJ)
• the shorter the covalent bond, the stronger the
  bond
• Br-F (237 kJ) gt Br-Cl (218 kJ) gt Br-Br (193 kJ)
• bonds get weaker down the column

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Using Bond Energies to Estimate DHrxn

• the actual bond energy depends on the surrounding
  atoms and other factors
• we often use average bond energies to estimate
  the DHrxn
• works best when all reactants and products in gas
  state
• bond breaking is endothermic, DH(breaking)
• bond making is exothermic, DH(making) -
• DHrxn ? (DH(bonds broken)) ? (DH(bonds
  formed))

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Estimate the Enthalpy of the Following Reaction
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Estimate the Enthalpy of the Following Reaction

• H2(g) O2(g) H2O2(g)
• reaction involves breaking 1mol H-H and 1 mol
  OO and making 2 mol H-O and 1 mol O-O
• bonds broken (energy cost)
• (436 kJ) (498 kJ) 934 kJ
• bonds made (energy release)
• 2(464 kJ) (142 kJ) -1070
• DHrxn (934 kJ) (-1070. kJ) -136 kJ
• (Appendix DHf -136.3 kJ/mol)

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Bond Lengths

• the distance between the nuclei of bonded atoms
  is called the bond length
• because the actual bond length depends on the
  other atoms around the bond we often use the
  average bond length
• averaged for similar bonds from many compounds

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Trends in Bond Lengths

• the more electrons two atoms share, the shorter
  the covalent bond
• CC (120 pm) lt CC (134 pm) lt C-C (154 pm)
• CN (116 pm) lt CN (128 pm) lt C-N (147 pm)
• decreases from left to right across period
• C-C (154 pm) gt C-N (147 pm) gt C-O (143 pm)
• increases down the column
• F-F (144 pm) gt Cl-Cl (198 pm) gt Br-Br (228 pm)
• in general, as bonds get longer, they also get
  weaker

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Bond Lengths
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Metallic Bonds

• low ionization energy of metals allows them to
  lose electrons easily
• the simplest theory of metallic bonding involves
  the metals atoms releasing their valence
  electrons to be shared by all to atoms/ions in
  the metal
• an organization of metal cation islands in a sea
  of electrons
• electrons delocalized throughout the metal
  structure
• bonding results from attraction of cation for the
  delocalized electrons

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Metallic Bonding
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Metallic BondingModel vs. Reality

• metallic solids conduct electricity
• because the free electrons are mobile, it allows
  the electrons to move through the metallic
  crystal and conduct electricity
• as temperature increases, electrical conductivity
  decreases
• heating causes the metal ions to vibrate faster,
  making it harder for electrons to make their way
  through the crystal

90
Metallic BondingModel vs. Reality

• metallic solids conduct heat
• the movement of the small, light electrons
  through the solid can transfer kinetic energy
  quicker than larger particles
• metallic solids reflect light
• the mobile electrons on the surface absorb the
  outside light and then emit it at the same
  frequency

91
Metallic BondingModel vs. Reality

• metallic solids are malleable and ductile
• because the free electrons are mobile, the
  direction of the attractive force between the
  metal cation and free electrons is adjustable
• this allows the position of the metal cation
  islands to move around in the sea of electrons
  without breaking the attractions and the crystal
  structure

92
Metallic BondingModel vs. Reality

• metals generally have high melting points and
  boiling points
• all but Hg are solids at room temperature
• the attractions of the metal cations for the free
  electrons is strong and hard to overcome
• melting points generally increase to right across
  period
• the charge on the metal cation increases across
  the period, causing stronger attractions
• melting points generally decrease down column
• the cations get larger down the column, resulting
  in a larger distance from the nucleus to the free
  electrons