Chapter 6

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Chapter 6 Chemical Bonding
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6-1 Introduction to Chemical Bonding

• Remember, the electrons in the outermost energy
  level, the valence electrons are involved in
  bonding.
• Atoms like to have a filled outer energy level
  (valence shell).
• Atoms can gain electrons to fill their valence
  shell.
• Atoms can lose electrons to reveal a full valence
  shell.
• Atoms can share electrons to get a filled valence
  shell.

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6-1 Chemical Bond

• A mutual electrical attraction between the nuclei
  and valence electrons of different atoms that
  binds the atoms together
• Most substances on earth are compounds all
  compounds contain chemical bonds they hold the
  atoms together

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6-1 Why do chemical bonds form?

• Most atoms by themselves have high potential
  energy
• They form bonds to minimize their potential
  energy
• Most atoms are less stable by themselves than
  combined with other atoms

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6-1 Types of Chemical Bonds

• Ionic bonding electrical attraction between a
  large number of cations and anions formation of
  ions involves transfer of electrons
• Covalent bonding sharing of electron pairs
  between atoms to form molecules
• Metallic bonding delocalized electrons give
  metals their special properties

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6-1 Ionic or Covalent?

• Bonding between two atoms is not usually purely
  ionic or purely covalent
• Remember, electronegativity is a measure of an
  atoms ability to attract electrons toward itself
  in a chemical bond
• Percent ionic/covalent character can be
  calculated using difference in electronegativity
  between two atoms

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6-1 Ionic or Covalent?

• Find the difference in electronegativity between
  the two elements in the bond
• The difference (absolute value) corresponds to
  the percentage ionic character
• Chart on right is in textbook p. 162

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6-1 Ionic or Covalent?

• gt50 - ionic
• 50 or less covalent
• Bonding between two identical atoms is completely
  covalent there is NO electronegativity
  difference between atoms of the same element
  the electrons are shared evenly
• Called NONPOLAR COVALENT BOND

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6-1 The Diatomic Molecules

• Diatomic molecule a molecule made of two atoms
  of the same element
• Seven elements exist naturally as diatomic
  molecules H, N, O, F, Cl, Br, I

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6-1 Polar Covalent Bonds

• Most covalent bonds are not purely covalent
• Most covalent bonds are polar covalent the
  electron density is shifted toward the more
  electronegative atom (uneven distribution of
  charge)
• Bonds with between 5 and 50 ionic character are
  POLAR COVALENT

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6-1 Polar Covalent, Nonpolar Covalent or Ionic?

• H-S
• Cs-S
• Cl-S
• Ca-Cl
• O-Cl
• Br-Cl

element electronegativity
H 2.1
S 2.5
Cs 0.7
Cl 3.0
Ca 1.0
O 3.5
Br 2.8
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6-2 Covalent Bonding and Molecular Compounds

• Molecule a neutral group of atoms that are held
  together by covalent bonds
• Molecular formula shows the types and numbers
  of atoms combined in a single molecule of a
  molecular compound
• BF3, CH3OH, CCl4

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6-2 Formation of a Covalent Bond

• Atoms far apart, dont influence each other
• Atoms approach each other, charged particles
  begin to interact (attractions and repulsions)
• Attractive force dominates until point 3, when
  attraction equals repulsion
• If atoms approach further, repulsion becomes
  increasingly greater, PE increases sharply

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6-2 Attractions and Repulsions Between Atoms
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6-2 Characteristics of the Covalent Bond

• BOND LENGTH distance between two bonded atoms
  at their minimum potential energy
• BOND LENGTH REPRESENTS A POTENTIAL ENERGY WELL!
• Forming a bond ALWAYS releases energy the same
  amount of energy must be ADDED to break the bond
  called BOND ENERGY

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6-2 Hydrogen Atoms in H2 Have Noble Gas
Configuration
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6-2 The Octet Rule

• Chemical compounds tend to form so that each
  atom, by gaining, losing or sharing electrons,
  has an octet of electrons in its highest occupied
  energy level
• F2, HCl, CH4
• Most main group elements form covalent bonds
  according to the octet rule

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6-2 Exceptions to the Octet Rule

• Incomplete octet BF3
• Expanded octet SF6

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6-2 Electron Dot Diagrams

• Electron configuration notation in which only the
  valence electrons are shown, indicated by dots
  placed around the elements symbol

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6-2 Lewis Structures

• Formulas in which atomic symbols represent nuclei
  and core electrons, dot-pairs or lines represent
  electron pairs (unshared electron pairs, bonding
  electron pairs)
• H2, F2, HF, NH3, CH4, CH3I, SH2

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6-2 Multiple Covalent Bonds

• Two atoms can share one, two or three pairs of
  electrons between them
• Double bond two pairs of electrons shared
  between atoms (O2)
• Triple bond three pairs of electrons shared
  between atoms (N2)

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6-2 Multiple Covalent Bonds

• Double bonds generally have HIGHER bond energies
  and SHORTER bond lengths than single bonds
• Triple bonds generally have HIGHER bond energies
  and SHORTER bond lengths than double bonds

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6-2 Sample Problems

• C2H6
• C2H4
• C2H2

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6-2 Resonance Structures

• Some molecules or ions cannot be adequately
  represented by one Lewis structure
• Resonance structures for ozone, O3

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6-2 Covalent-Network Bonding

• Do not contain individual molecules
• Continuous, three dimensional networks of bonded
  atoms
• Ex. graphite

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Diamond A Covalent Network
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6-3 Ionic Bonding and Ionic Compounds

• Ionic compound composed of positive and
  negative ions that are combined so that the
  numbers of positive and negative charges are
  equal
• Formula unit simplest collection of atoms from
  which an ionic compounds formula can be
  established simplest whole number ratio of
  cations to anions that will give a neutral formula

• calcium fluoride
• sodium oxide
• magnesium sulfide

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6-3 Formation of Ionic Bonds (represented with
dot diagrams)

• sodium chloride
• calcium fluoride
• magnesium sulfide

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6-3 Characteristics of Ionic Bonding

• Ions arrange themselves to minimize potential
  energy
• Oppositely charged ions attract each other
• Cations surrounded by anions and vice versa
• Arrangement is called a crystal lattice

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6-3 Strength of Ionic Bonds

• Bond formation releases energy
• Lattice energy energy released when one mole of
  an ionic crystalline compound is formed from
  gaseous ions
• Negative values indicate energy is released
• More negative stronger bond
• Table 6-3 (p. 179)

compound Lattice energy (kJ/mol)
NaCl -787.5
NaBr -751.4
CaF2 -2634.7
CaO -3385
LiCl -861.3
LiF -1032
MgO -3760
KCl -715
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6-3 Ionic vs. Molecular Compounds

• Forces that hold ions together are very strong
• Covalent bonds also very strong, but forces of
  attraction between molecules (intermolecular
  forces) much weaker

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6-3 Ionic v. Molecular Compounds
Ionic Molecular
Forces
Melting Point
Boiling Point
Hardness
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6-3 Why are ionic compounds brittle?

• Shifting ions slightly puts like charges next to
  each other

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6-3 Solubility of Ionic Compounds

• Polar water molecules pull ions away from the
  crystal and surround them.
• Many ionic compounds are soluble in water.

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6-3 Solubility of Ionic Compounds

• In solid state, ions cant move, cant conduct
  electricity
• When ionic compounds dissolve in water, the
  charged particles are free to move the solution
  can conduct electricity

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6-3 Polyatomic Ions

• Monatomic ions form when a single atom gains or
  loses an electron or electrons
• Polyatomic ions form when a group of atoms that
  are bonded covalently take on a charge

• ammonium
• nitrate
• sulfate
• carbonate

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6-3 Lewis Structures of Polyatomic Ions/Resonance

• Ammonium
• Nitrate

• Sulfate
• Carbonate

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6-4 Metallic Bonding

• The unique properties of metals can be accounted
  for by the metallic bond.
• Conduct heat and electricity, malleable, ductile,
  luster

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6-4 Electron Sea Model

• Metals have only 1, 2 or 3 valence electrons
• Also have vacant p- and d- orbitals
• When metal atoms are close to each other, these
  vacant orbitals overlap
• Outer electrons roam freely throughout network of
  overlapping orbitals electrons are delocalized
• Metallic bonding results from attraction
  between metal atoms and the surrounding sea of
  electrons

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6-4 Conductivity and Luster

• When charged particles are free to move, an
  electrical current can pass through metals
  conduct electricity
• Because metal atoms have many orbitals separated
  by small energy differences, metals absorb many
  light frequencies when energy is emitted, light
  is released looks shiny

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6-4 Malleability and Ductility

• When atoms are moved, electrons flow around them
  and take new shape

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6-4 Metallic Bond Strength

• Varies with nuclear charge of atoms and number of
  electrons in electron sea (metals with one
  valence electron are softer than metals with 2
  valence electrons)
• Heat of vaporization heat required to vaporize
  a metal is a measure of the strength of the bonds
  that hold it together

period Heats of Vaporization, kJ/mol Heats of Vaporization, kJ/mol Heats of Vaporization, kJ/mol
2nd Li, 147 Be, 297
3rd Na, 97 Mg, 128 Al, 294
4th K, 77 Ca, 155 Sc, 333
5th Rb, 76 Sr, 137 Y, 365
6th Cs, 64 Ba, 140 La, 402
Table 6-4 on p. 182
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6-5 Molecular Geometry

• VSEPR Valence Shell Electron Pair Repulsion
• VSEPR theory is a model that accounts for the
  shapes of simple molecules

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6-5 VSEPR Theory

• Repulsion between sets of valence electrons
  surrounding an atom causes these sets to be
  oriented as far apart as possible
• BeF2

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6-5 VSEPR Theory

• BF3
• CH4

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6-5 VSEPR

• NH3
• H2O

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A central atom X atom bonded to central
atom E unshared electron pair on central atom

• The shapes of simple molecules are determined by
  the number of atoms bonded to the central atom
  and the number of unshared pairs of electrons
  around the central atom.

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6-5 VSEPR

• SF2
• PCl3

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6-5 VSEPR

• CHCl3
• CO2

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6-5 Hybridization

• Hybridization is a model that explains how the
  orbitals of an atom are rearranged when the atom
  forms covalent bonds
• Hybridization is the mixing of two or more atomic
  orbitals of similar energies on the same atom to
  produce new orbitals of equal energies
• Especially useful for explaining bonding in
  carbon compounds

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6-5 Methane

• Methane has tetrahedral geometry (predicted by
  VSEPR and known from experimentation), but
  valence electrons of carbon atom are in 2
  different kinds of orbitals
• How does carbon make four equivalent covalent
  bonds in this compound?

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6-5 Methane

• ___ ___ ___
• 2p
• ___ ___ ___ ___
• sp3
• ___
• 2s

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6-5 Methane
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6-5 Hybrid Orbitals

• Orbitals of equal energy produced by the
  combination of two or more orbitals on the same
  atom
• Number of hybrid orbitals equals number of atomic
  orbitals that have combined

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6-5 Hybridization

Atomic Orbitals Type of Hybridization Number of Hybrid Orbitals
s, p sp 2
s, p, p sp2 3
s, p, p, p sp3 4
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6-5 Ethane
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6-5 Ethene
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6-5 Ethyne
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6-5 Intermolecular Forces (van der Waals Forces)

• Forces of attraction between molecules
• Vary in strength but generally weaker than ionic,
  metallic or covalent bonds
• Boiling point is a good measure of the strength
  of intermolecular forces

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bonding type substance bp (1 atm, C)
nonpolar-covalent H2 -253
(molecular) O2 -183
Cl2 -34
Br2 59
CH4 -164
CCl4 77
C6H6 80
polar-covalent PH3 -88
(molecular) NH3 -33
H2S -61
H2O 100
HF 20
HCl -85
ICl 97
ionic NaCl 1413
MgF2 2239
metallic Cu 2567
Fe 2750
W 5660

• Table 6-7, p. 190

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6-5 Molecular Polarity and Dipole-Dipole Forces

• Polar molecules (like water) are dipoles. They
  have two poles, one positive and one negative.
• Forces of attraction between polar molecules are
  called dipole-dipole forces.

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6-5 Dipole-Dipole Forces

• Short range
• Act only between nearby molecules
• Polarity of molecules is determined by types of
  bonds and arrangement of bond

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6-5 Polarity

• Water
• ammonia

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6-5 Polarity

• carbon tetrachloride
• carbon dioxide

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6-5 Dipole-Induced Dipole

• Electron clouds are mobile
• A permanent dipole (like water, ammonia, hydrogen
  chloride) can induce a temporary dipole in a
  nonpolar molecule

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6-5 Hydrogen Bonding

• Occurs in compounds in which hydrogen is attached
  to oxygen, nitrogen or fluorine
• Very strong
• Many of waters special properties can be
  accounted for by hydrogen bonding

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6-5 London Dispersion Forces

• Weak forces
• Electrons are in constant motion
• Molecules can have temporary dipoles due to this
  movement of electrons
• A temporary dipole can induce another temporary
  dipole

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• Fatty acids can be saturated of unsaturated.
• Saturated fatty acids have all single bonds.
• Unsaturated fatty acids have some double bonds.
• Double bonds cause a kink in the carbon chain.
• Unsaturated fatty acids dont pack together as
  well, have weaker dispersion forces between them,
  are less likely to form solid in arteries.

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