CHAPTER 9: BONDING AND MOLECULAR STRUCTURE

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CHAPTER 9 BONDING AND MOLECULAR STRUCTURE

• "I have attempted to give you a glimpse...of what
  there may be of soul in chemistry.
• -G.N. Lewis

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9.0 Objectives

• Understand the basic process of ionic bonding
  identify ionic compounds and describe their
  internal structure and properties, understand how
  size of ions affects ionic properties, calculate
  lattice energy, and draw Lewis diagrams of ionic
  structures.
• Identify covalent compounds and characterize
  their properties.
• Draw Lewis structures of covalent substances
  including exceptions such as reduced and expanded
  octets, radicals, and resonance structures.
• Define and predict trends in bond order, bond
  length, and bond dissociation energy. Use bond
  energy to predict enthalpy of a reaction.
• Understand the concept of electronegativity and
  how it is used to predict polarity of individual
  bonds and entire molecules.
• Use VSEPR Theory to predict the shapes of simple
  covalent molecules.

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Homework

• HW1 29, 33, 38, 39, 40, 41, 43, 45
• Valence e-, LEWIS STRUCTURES!
• HW2 47, 49, 51, 53, 95
• Formal Charge, Polarity, Electonegativity
• HW3 69, 93, 109
• Bond Energy
• HW4 73, 75, 77, 79, 81, 89, 99, 103
• Molecular Geometry
• HW5 83, 85, 97
• Polarity

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9.1 VALENCE ELECTRONS

• 1. Bonding definition
• forces that hold atoms together
• What role do the e- play?
• 2. Valence electrons vs. Core electrons
• MAIN GROUP Elements
• Outermost s and p e-
• Transition Metals
• Outermost s and p e- as well as (n-1) d e-
• When in doubt, write the noble gas e-
  configuration

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9.1 VALENCE ELECTRONS

• 3. Lewis dot diagrams of elements
• Diagrams that showcase valence e-
• Lewis says, Place the first four dots
  separately!
• Ex. Li, Be, B, C, N, O, F, Ne

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9.2 CHEMICAL BOND FORMATION

• 1. Ionic bonding definition and Lewis
  representation
• Bond between metal and nonmetal due to
  electrostatic interactions
• Metal donates e-
• Nonmetal accepts e-
• Ex. NaCl and Na2S

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9.2 CHEMICAL BOND FORMATION

• 2. Covalent bonding and Lewis representation of
  Cl2
• Bond in which e- are shared
• Overlap of e- density between 2 orbitals
• Ex. Cl2

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9.2 CHEMICAL BOND FORMATION

• 3. Continuum
• Complete ionic or complete sharing of e- is a bit
  extreme most bonding has uneven sharing of e-
  (sometimes ionic, sometimes covalent)
• 4. Other bond types
• Metallic bonding
• Ex. Alloys

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9.3 BONDING IN IONIC COMPOUNDS

• 1. Steps in formation of NaCl
• 1. Na(g) ? Na(g) e- ?E 496
  kJ/mol
• 2. Cl(g) e- ? Cl-(g) ?E -349
  kJ/mol
• 3. Na(g) Cl-(g) ? Na, Cl- ?E -498 kJ/mol
• ?Eoverall -351 kJ/mol

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9.3 BONDING IN IONIC COMPOUNDS

• 2. Lattice energy
• energy for the formation of 1 mol of solid
  crystalline ionic compound when ions in the gas
  phase combine

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9.3 BONDING IN IONIC COMPOUNDS

• 3. Formula units
• RECALL Smallest repeating unit of an ionic
  compound

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9.4 COVALENT BONDING AND LEWIS STRUCTURES

• 1. Diagram of H2
• Single Hydrogens
• H? ?H Both want 1s2
• 1 shared pair
• HH
• Bonds are represented as single lines
• HH

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9.4 COVALENT BONDING AND LEWIS STRUCTURES

• 2. Orbital overlap diagrams of H2, HCl, Cl2

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9.4 COVALENT BONDING AND LEWIS STRUCTURES

• 3. Terminology single, double, and triple
  bonds, bonding pairs and nonbonding or lone pairs
  of electrons
• Single Bond 2 e- shared between 2 atoms
• Ex. H2
• Double Bond 4 e- shared between 2 atoms
• Ex. O2
• Triple Bond 6 e- shared between 2 atoms
• Ex. N2

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9.4 COVALENT BONDING AND LEWIS STRUCTURES

• Bonding Pairs e- involved in bonding
• (See preceding examples)
• Nonbonding (lone) pairs e- that are not involved
  with bonding but help provide the octet for an
  atom
• Ex. Cl2

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9.4 COVALENT BONDING AND LEWIS STRUCTURES

• 4. Octet Rule
• The INFAMOUS noble-gas configuration
• tendency for molecules/polyatomic ions to have
  structures in which 8 e- surround each atom
• H, He have a duet

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9.4 COVALENT BONDING AND LEWIS STRUCTURES

• 5. Rules for drawing Lewis structures
• a. Choose a central atom
• Usually the atom with the lowest e- affinity
• Usually makes a lot of bonds
• Halogens are generally terminal atoms
• b. Count the total number of valence electrons
• Neutral Molecule sum of valence e- for each atom
• Anions sum of valence e- and negative charge
• Cations valence e- minus the total positive
  charge

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9.4 COVALENT BONDING AND LEWIS STRUCTURES

• c. Draw a skeleton structure
• Use one pair of electrons to form a bond between
  each pair of bound atoms
• d. Place the remaining electrons to fulfill the
  octet rule
• Do this for each atom
• Hydrogen gets a duet

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9.4 COVALENT BONDING AND LEWIS STRUCTURES

• e. Lack of electrons
• Requires multiple bonds (double, triple)
• Could be more than one multiple bond
• f. Too many electrons
• Verify that your structure is correct (octets for
  all?)
• Watch anions!

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9.6 Lewis Structures of Some Simple Molecules

• O-VS
• S Shared e- in bonds
• O total e- required for an Octet
• V Valence e- for all elements

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9.4 COVALENT BONDING AND LEWIS STRUCTURES

• 6. Diagrams of H2 F2 CH4 NH3 H2O HF OH-
  NH4

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9.4 COVALENT BONDING AND LEWIS STRUCTURES

• H2 F2 CH4 NH3 H2O HF OH- NH4

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9.4 COVALENT BONDING AND LEWIS STRUCTURES

• 7. Isoelectronic species NO N2 CO CN-

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9.5 RESONANCE

• 1. Definition
• Alternative and equivalent Lewis structure
  created by shifting the e- in a structure
• Spinning Rim Analogy

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9.5 RESONANCE

• 2. Examples NO3- and NO2-

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9.5 RESONANCE

• 3. Experimental evidence says
• Its a combination of both
• There are however, MORE PREVALENT resonance
  structures for some molecules
• Benzene is the most classic of all resonance
  structures

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9.5 RESONANCE
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9.6 EXCEPTIONS TO THE OCTET RULE

• 1. Reduced octets for H, B and Be
• Ex. BeCl2, BCl3 (Be 4 e-, B 6e-)

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9.6 EXCEPTIONS TO THE OCTET RULE

• 2. Expanded octets PF5 SF6 ClF4- XeF2
• Watch these elements (and some others) for
  expanded octets P, S, Cl, As, Se, Br, Kr, Xe

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9.6 EXCEPTIONS TO THE OCTET RULE

• 3. Radicals (paramagnetic) NO and NO2
• Structure that has unpaired e-
• Extremely Reactive

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9.6 EXCEPTIONS TO THE OCTET RULE

• 3. Problems with Lewis structures
• Only show 2-D view? life (chemistry) is 3-D
• Works for most molecules, but not all
• Doesnt show how evenly/unevenly e- are being
  shared

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9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND
MOLECULES

• 1. Definitions polar and nonpolar bonds
• Nonpolar bonds 2 e- in a bond are evenly
  shared between the 2 atoms
• Polar bonds 2 e- in a bond are unevenly shared
  one atom is taking more of the e- density atoms
  have a partial charge

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9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND
MOLECULES

• 2. Electronegativity
• a. definition
• (EN) ability of an atom to attract bonding e- to
  itself when the atom is in a molecule
• b. Table and Periodic trends
• See Pg.10 in Reference Booklet
• Increases going left to right and bottom to top
• (Fluorine greatest at 4.0)

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9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND
MOLECULES
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9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND
MOLECULES

• 3. ?EN parameters
• Prediction of Ionic Character
• Pure Covalent Pure Ionic
• 0 .5 1 1.5 2 2.5 3
• In General 0.0 lt 0.45 Nonpolar
• 0.45 1.8 Polar Covalent
• gt 1.8 Ionic

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9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND
MOLECULES

• 4. Ex9.1 Arrange the following bonds in order of
  increasing polarity F-Cl, F-F, F-Na

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9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND
MOLECULES

• 5. Central atom in Lewis structure
• Many times has a formal charge
• Making more/less bonds than it normally does

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9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND
MOLECULES

• 6. Formal Charge
• a. Definition and Use
• Charge for an atom in a molecule based on premise
  that bonding e- are evenly shared
• b. Calculating equation
• Formal Charge Group - Lone Pair e- ½
  Bonding e-

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9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND
MOLECULES

• c. Examples OH- and NO3-

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9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND
MOLECULES

• Ex9.2 Calculate the formal charge on each atom
  in CO32- and NH4

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9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND
MOLECULES

• 7. Electroneutrality Definition
• The e- in a molecule are distributed so that the
  formal charge is minimal
• Most Probable Lewis Structure one with minimal
  FC minimal FC is more important than symmetry
• Negative charge should reside on the most
  electronegative element
• Formal charge gt /- 2 is not likely

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9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND
MOLECULES

• a. Example CO2

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9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND
MOLECULES

• b. Ex9.3 Use formal charge and the
  electroneutrality concept to determine the most
  likely structures for N2O and OCN1-

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9.8 BOND PROPERTIES

• 1. Bond order
• a. definition and examples
• number of bonding e- pairs shared between 2 atoms
• Usually an integer (1, 2, or 3)
• BOND ORDER ( shared pairs linking
  X-Y)_____ (number of X-Y links
  in the molecule)
• Ex. CH4, CO2
• b. resonance structures
• Bond order are fractions
• e- residing over both locations evenly
• Ex. O3

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9.8 BOND PROPERTIES

• 2. Bond Length definition and examples
• Bond length distance between nuclei in a
  covalent bond
• More Polar bonds shorter length
• More bonds shorter length
• Ex. C-C CC CC
• 1.54Å 1.34Å 1.20Å

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9.8 BOND PROPERTIES

• 1. Bond dissociation energy definition and
  examples
• Bond Dissociation Energy (D)
• energy needed to break a covalent bond in the gas
  phase
• Higher Bond Order? Higher D

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9.8 BOND PROPERTIES

• a. Estimating Enthalpy of reaction from bond
  energies equation
• ?Hrxn ? D (bonds broken) - ? D (bonds formed)
• Energy is required to break bonds
• Energy is released when bonds are formed

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9.8 BOND PROPERTIES
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9.8 BOND PROPERTIES

• b. Example Estimate the ?HRxn for the
  synthesis reaction between gaseous hydrogen and
  chlorine.

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9.8 BOND PROPERTIES

• c. Ex9.4 Estimate the enthalpy of reaction for
  the combustion of methane, CH4, to produce
  gaseous carbon dioxide and water vapor.

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9.9 MOLECULAR SHAPES

• 1. Gumdrops and toothpicks
• A tasty way to practice
• chemistry!

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9.9 MOLECULAR SHAPES

• 2. VSEPR Theory and importance of shapes
• VSEPR Valence Shell Electron Pair Repulsion
• VSEPR gives our 2-D Lewis structures LIFE (3-D)
• Geometry has HUGE impact on properties
• Based on idea pairs of e- in bonded atoms repel
  one another
• Want to be as far apart as possible? Gives shape
• Electron Group
• Any collection of valence e- around an atom that
  repel other e-
• Single unpaired e-
• Lone pair e-
• Bonding pairs of e- (1, 2, 3)

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9.9 MOLECULAR SHAPES

• NOTATION
• AXnEm
• A central atom
• X terminal atoms
• E lone pair e- on central atom

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9.9 MOLECULAR SHAPES

• 3. Single bonds, no unshared pairs of electrons
• Lewis structure, geometry and bond angles of
• a. BeH2

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9.9 MOLECULAR SHAPES

• BH3
• CH4

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9.9 MOLECULAR SHAPES
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9.9 MOLECULAR SHAPES

• 4. Unshared pairs of electrons on the central
  atom
• a. NH3
• b. H2O

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9.9 MOLECULAR SHAPES

• c. analogs (H2S, PCl3, etc.)

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9.9 MOLECULAR SHAPES

• Ex9.5 Predict the molecular geometry and bond
  angles of HOCl and SiO44-

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9.9 MOLECULAR SHAPES

• 6. Multiple Bonds
• a. CO2
• b. H2CO

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9.9 MOLECULAR SHAPES

• c. HCN

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9.9 MOLECULAR SHAPES

• 7. Ex9.6 Predict the molecular geometry and bond
  angles in the following species C2H2 C2H4
  ClO31- NO31- N2O ONCl

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9.9 MOLECULAR SHAPES

• 7. Expanded octets
• a. definition and recognizing
• More than 4 e- groups around a central atom
• Use formal charge to help guide Lewis
  Structure/Geometry

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9.9 MOLECULAR SHAPES

• b. Examples
• PCl5
• SF6

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9.9 MOLECULAR SHAPES

• ClF5
• XeF4

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9.9 MOLECULAR SHAPES

• 8. Ex9.7 Determine the molecular geometry and
  bond angles of ICl2-1 IF3 XeOF4 How does the
  electronic geometry differ from the molecular
  geometry for these species?

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9.10 MOLECULAR POLARITY

• 1. Nonpolar molecules definition and examples,
  CH4 CO2
• Molecule consisting entirely of nonpolar bonds
  OR
• Molecule with polar bonds that cancel one another
  out

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9.10 MOLECULAR POLARITY

• 2. Dipoles definition and examples, NH3
  H2O
• Molecule with separate centers of () and (-)
  charge
• Polar bonds present with no canceling out

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9.10 MOLECULAR POLARITY

• 4. Dipole moments and vectors (?)
• ? ? d
• ? magnitude of charge
• d distance
• Measured in debye (D)
• Nonpolar ? 0
• Polar ? ? 0

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9.10 MOLECULAR POLARITY

• 5. Rules for determining polarity of a molecule.
  A molecule is a dipole if
• (uneven balance of e- density)
• RULES
• 0. Draw Lewis Structure
• 1. Use VSEPR to predict molecular shape
• 2. Electronegativity to predict bond dipoles
• 3. Determine whether bond dipoles cancel to
  produce nonpolar or combine to give polar
  molecule
• KEY CLUES
• 1. Polar Bonds
• 2. Lone Pair e-
• 3. 2 different atoms bonded to central atom

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9.10 MOLECULAR POLARITY

• 6. Ex9.8 Determine which of the following are
  dipoles SO2 BF3 CO2 N2O ClO3- ONCl
  NCl3 BFCl2 SCl2

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9.10 MOLECULAR POLARITY

• 6. Exceptions XeF4

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9.10 MOLECULAR POLARITY

• 7. Properties of dipoles
• Special properties observed due to interactions
  between molecules (intermolecular forces)
• Attraction to other dipoles

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END OF CHAPTER 9!