Chemical Bonding

1
Chemical Bonding
2
What is a Bond?

• Force that holds atoms together
• Results from the simultaneous attraction of
  electrons (-) to the nucleus ()

3
Breaking/Forming Bonds

• When a bond is broken energy is absorbed
• Endothermic
• When a bond is formed energy is released
• Exothermic
• The greater the energy released during the
  formation of the bond, the greater its stability
• Stable bonds require a great deal of energy to
  break

4
Lewis Dot Diagrams

• Use dots to represent the number of valence
  electrons
• How to write
• Write the symbol.
• Put one dot for each valence electron
• Electrons go on the 4 sides, no more than 2 per
  side

5
Dot Diagram Examples

• Draw dot-diagrams for the following
• Mg
• C
• Ne

6
Dot Diagrams - Ions

• For ions, use brackets and place the charge
  outside the brackets
• Examples
• Na
• O2-
• H

7
Octet Rule

• Atoms will gain or lose electrons in order to
  have a full valence shell like the nobles gases
• Take the shortest route
• Metals lose electrons to form positive ions
  (Cations)
• Nonmetals gain electrons to form negative ions
  (Anions)

8
Exceptions

• 1st principle energy level only holds 2 electrons
• Transition elements can lose valence (s) and
  inner (d) electrons this is why they have
  multiple oxidation states
• Some atoms may be stable with less than an octet
  many compounds with B
• Some atoms may be stable with more than an octet
  elements beyond period 2, especially P and S,
  the additional electrons are added to the d
  sublevel
• Molecules with an odd number of electrons they
  will be unstable

9
Types of Bonds

• Ionic - Electrons are transferred from a metal to
  a nonmetal
• Covalent - Electrons are shared between 2
  nonmetals
• Polar Covalent electrons are shared unequally
• Nonpolar Covalent electrons are shared equally
• Metallic - Electrons are mobile within a metal,
  Sea of Electrons

10
Dog Analogy

• Ionic Bonds
• big greedy dog stealing the other dogs bone
• Polar Covalent Bonds
• Unevenly matched but willing to share
• Nonpolar Covalent Bonds
• Dogs of equal strength
• Metallic Bonds
• Mellow dogs with plenty of bones to go around
• See the Dogs

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Identifying Bond Type

• Ionic metal and a nonmetal
• Covalent 2 nonmetals
• Metallic metals
• OR Use electronegativity differences
• Ionic 1.7 or more
• Polar Covalent 0.5-1.6
• Nonpolar Covalent 0.0-0.4

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Identifying Bond Types

• Indicate the type of bond present in each
• HCl
• CCl4
• MgCl2
• O2
• Hg
• H2O

13
Ionic Bonds

• Transfer of 1 or more electrons from a metal to a
  nonmetal
• Electronegativity difference is 1.7
• Example Sodium Chloride (NaCl)

Na electron transferred to Cl
Na Cl
X
14
Monatomic Ions

• One atom in an ion
• Look at the valence electrons to determine the
  charges
• Examples K, O2-

15
Polyatomic Ions

• More than one atom in the ion
• Reference Table E
• Charge belongs to the entire ion, not an
  individual atom
• Within the polyatomic ion the atoms are held
  together by covalent bonds
• When writing it, place ( ) around the entire ion,
  with the charge outside
• Examples (NH4), (H3O), (CO3)2-

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Writing Ionic Formulas

• You need an equal amount of positive and negative
  charges, so that the compound is neutral
• Ionic Formulas are always written as empirical
  formulas (reduced)

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Examples

• Na1 Cl1-
• Mg2 Cl1-
• Ca2 CO32-
• Al3 O2-

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Criss Cross Method

• Write the symbol for the cation and anion
• Write each ions charge as a superscript
• Criss-cross the charges to become subscripts of
  the other ion
• Do not put () or (-) charges in the final
  formula
• Reduce to least common multiple (empirical
  formula)

19
Ionic Formulas

• Write the formula for the compound formed from
  the following ions
• Mg2 Cl-
• Ca2 CO32-
• Al3 O2-
• Ca OH

20
Naming Ionic Compounds

• Name the cation first, the anion second
• Cation keeps its name, anion changes its ending
  to ide (Chlorine ? Chloride)
• Do not change the ending of polyatomic ions
• Examples
• NaCl
• CaCO3
• MgF2

21
Stock System only used for positive ions

• Some cations have more than one positive
  oxidation states
• A roman numeral is used to indicate the charge of
  the positive ion

22
Stock System Examples

• Iron (II) Chloride
• Iron (III) Oxide
• Copper (II) Oxide
• a. What charge does copper have in copper II
  sulfate?
• b. What is the formula for copper II sulfate?

23
Ionic Salts

• Salts are ionic compounds made up of cations and
  anions
• The ratio of cations to anions is always such
  that an ionic compound has no overall charge
• Many of the ions are bonded together to form a
  crystal

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Properties of Ionic Salts

• Ionic Bonds are very strong
• Very high melting and boiling points
• Hard
• Brittle

25
Properties of Salts (contd)

• Do not conduct electricity as solids
• Do conduct electricity when the salt melts or is
  dissolved in water (liquid phase or aqueous)
• In order to conduct electricity a substance must
  have free moving charged particles
• In the solid phase the ions are not free to move

26
Melting and Boiling Points of Compounds
27
Covalent Bonds

• Sharing of electrons between 2 nonmetals
• Electronegativity difference is 1.6

28
Non-Polar Covalent

• Electrons are shared equally
• Uniform distribution of electrons
• Bond is symmetrical
• Electronegativity difference of 0-0.4
• All diatomic molecules have non-polar covalent
  bonds

29
Nonpolar Covalent Examples

• Flourine (F2)
• e-neg difference
• Dot diagram
• Hydrogen (H2)
• e-neg difference
• Dot diagram

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Polar Covalent

• Unequal Sharing of electrons
• Unequal distribution of electrons
• Partial positive and partial negative charges
• The side with the higher electronegativity will
  have a greater share of the electron(s) resulting
  in a partial negative charge
• Electronegativity difference of 0.5-1.6

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Polar Covalent Examples

• HCl
• e-neg difference
• Dot diagram
• H2O
• e-neg difference
• Dot diagram

32
Dipoles

• Form when the charge in a bond is asymmetrical
• Present in polar bonds
• Partial positive and partial negative charges

33
Polar Bonds / Dipoles

• Isnt a whole charge just a partial charge
• means a partially positive
• means a partially negative
• Example
• H - Cl
• ---?
• The Cl pulls harder on the electrons (more eneg)
• The electrons spend more time near the Cl

d
d-
d
d-
34
Dipole Examples

• Which molecule contains more polar bonds?
• a. CCl4
• b. CH4
• 2. Which has a stronger dipole?
• HCl
• HBr

35
Properties of Molecular Substances (Covalent
Compounds)

• Soft
• Low melting points and boiling points
• Many exist as gases
• Poor conductors of heat and electricity (in all
  phases)
• Examples H2O, CCl4, NH3, C6H12O6, O2

36
Molecular Formulas (Covalent Compounds)

• Contain covalent bonds
• Tells you how many atoms are present in a single
  molecule
• Named similarly to ionic compounds, except use
  prefixes to indicate the number of atoms per
  molecule

37
Prefixes

• Mono- is only used for the second element
• Example CO carbon monoxide

38
Examples

• CCl4
• H2O
• NO
• N2O5
• BBr3

39
Structural Formulas

• Specifies how atoms are bonded together
• Dashes represent bonds
• 2 atoms can share up to 3 pairs of electrons

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Single Bonds

• 2 atoms share 1 pair of electrons (2 electrons)
• Examples
• Ammonia (NH3)
• Chlorine (Cl2)
• Hydrochloric Acid (HCl)

41
Double Covalent Bonds

• 2 atoms share 2 pairs of electrons (4 electrons)
• 2 bonds between 2 atoms
• Examples
• Carbon Dioxide (CO2)
• Oxygen (O2)

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Triple Covalent Bond

• 2 atoms share 3 pairs of electrons (6 electrons)
• 3 bonds between 2 atoms
• Examples
• Nitrogen (N2)
• Ethyne (C2H2)

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Bond Length/Strength

• Length
• Single gt Double gt Triple
• The more electrons in a bond, the greater the
  attraction, therefore shorter
• As you move down a group bond length increases
• Due to increasing molecular size
• Strength
• Triple is the strongest, most stable, requires
  the most energy to break

44
Network Solids

• Covalently bonded atoms are linked into a giant
  network (macromolecules)
• Examples Diamond (C), Graphite (C), Silicon
  Carbide (SiC), and Silicon Dioxide (SiO2)

45
Network Solids

• Properties
• Hard
• High melting and boiling points
• Do not conduct heat and electricity

46
Metallic Bonding

• Sea of Electrons
• Electrons are free to move through the solid.

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Properties of Metallic Solids

• Very Strong
• Good conductors of heat and electricity because
  electrons are free to move about
• Luster
• High melting point (except Hg)
• Malleable, Ductile

48
VSEPR Theory

• In a small molecule, the electron pairs are as
  far away from each other as possible
• VSEPR Valence Shell Electron Pair Repulsion

49
Linear

• Drawn on a straight line
• All molecules of only 2 atoms are linear
• Many 3 atom molecules are linear, if there are no
  unshared electron pairs on the central atom
• If both ends are the same, the molecule is
  nonpolar (Symmetrical Nonpolar)
• If the ends are different, the molecule will be
  polar (Asymmetrical Polar)
• Bond Angle 180o
• See Molecules
• Examples H2, CO2, HCl

50
Trigonal Planar
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Trigonal Planar

• A central atom is bonded to 3 other atoms, with
  no extra electrons on the central atom
• Forms a flat Y shape (triangle shape)
• If the ends are all the same, NONPOLAR
• If the ends are different, POLAR
• Bond Angle 120o
• See Molecules
• Examples BCl3, BH2F

52
Pyramidial

• A central atom is bonded to 3 other atoms and the
  central atom has an unshared electron pair
• 3-D, like a pyramid
• Always POLAR
• Bond Angle 107o
• See Molecules
• Example NH3

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Tetrahedral
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Tetrahedral

• A central atom bonded to 4 other atoms
• 3-D shape allows the electron pairs to get as far
  away from each other as possible

H
109.5º
C
H
H
H
55
Tetrahedral

• If all the ends are the same, NONPOLAR
• If the ends are different, POLAR
• Bond Angle 109.5o
• See Molecules
• Examples
• 1. CH4
• 2. CH3Cl

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Bent
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Bent

• A central atom is bonded to 2 other atoms and the
  central atom has 2 unshared electron pairs
• Always POLAR
• Bond angle 105o
• See Molecules
• Example H2O

58
Intermolecular Attractions/Forces

• Forces between molecules
• Determines boiling point, melting point, vapor
  pressure, surface tension
• The stronger the intermolecular attractions, the
  higher the boiling point
• All intermolecular attractions are weaker than
  actual bonds

59
Dipole-Dipole Forces

• Occurs between 2 polar molecules
• The positive end of one molecule is attracted to
  the negative end of another molecule
• The greater the electronegativity difference is,
  the more polar the bond will be and the stronger
  the dipole will be
• Example HCl

60
Dipole Examples

• 1. Which would have the strongest intermolecular
  forces? Explain Why.
• a. HCl
• b. HBr
• 2. Which would have the weakest intermolecular
  forces? Explain Why.
• a. H2S
• b. H2O

61
Hydrogen Bonds

• Special, Strong type of Dipole Attractions
• Attraction of a covalently bonded H atom to a F,
  O, or N atom on another covalent compound

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Hydrogen Bonds

• VERY STRONG
• Molecules with H bonds will have high boiling
  points, melting points, and surface tension
• Example NH3

63
H-bonds Examples

• 1. Which sample has Hydrogen Bonds?
• a. H2 b. HF c. F2 d. HCl
• 2. Which is the strongest?
• a. Hydrogen Bonds
• b. Covalent Bonds
• c. Dipole-Dipole Attractions

64
Molecule Ion Attractions

• Attraction between a polar compound and an ion
  (ionic salt)
• Polar substances (such as water) attract ions
  from ionic compounds in solution
• This allows ionic substances to dissolve in polar
  solvents (water)
• The anion is attracted to the positive end of the
  polar solvent
• The cation is attracted to the negative end of
  the polar solvent
• The ion dissociates (falls apart)
• Example NaCl(aq)

65
Molecule-Ion Examples

• 1. Molecule-Ion attractions are present in which
  sample?
• a. HCl(l) c. KCl(l)
• b. HCl(aq) d. KCl(aq)
• 2. When sodium chloride dissolves in water the
  chloride ion is attracted to
• a. The positive part of the water, the O atom
• b. The negative part of the water, the O atom
• c. The positive part of the water, the H atom
• d. The negative part of the water, the H atom

66
Van Deer Waals Forces

• Very weak
• Exist between non-polar molecules
• Caused by momentary dipoles
• Increases as molecular mass increases

67
VDW Examples

• 1. Rank in order from weakest to strongest
• Hydrogen Bonds
• Covalent Bonds
• Van deer Waals Forces
• Dipole-Dipole Attractions
• 2. Which would have the strongest intermolecular
  forces?
• a. H2 b. Cl2 c. F2 d. Br2