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
Identifying Bond Type

• Ionic metal and a nonmetal
• Covalent 2 nonmetals
• Nonpolar Covalent (equal)
• Same atoms (diatomics, triatomics)
• Polar Covalent (unequal)
• Unequal electronegativity
• Metallic metals

11
Identifying Bond Types

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

12
Ionic Bonds

• Transfer of 1 or more electrons from a metal to a
  nonmetal
• The greater the electronegativity difference
  between atoms, the greater the ionic character
• Example Sodium Chloride (NaCl)

Na electron transferred to Cl
Na Cl
X
13
Monatomic Ions

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

14
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-

15
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

1. Na1 Cl1-
2. Mg2 Cl1-
3. Ca2 CO32-
4. Al3 O2-

17
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)

18
Ionic Formulas

• Write the formula for the compound formed from
  the following ions
• Mg2 Cl-
• Ca2 CO32-
• Al3 O2-
• Calcium ion hydroxide ion

19
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

20
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

21
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?

22
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

23
Properties of Ionic Salts

• Ionic Bonds are very strong
• Very high melting and boiling points
• All solids at STP
• Hard
• Brittle

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Melting and Boiling Points of Compounds
Compound Name Formula Type of Compound mp (oC) bp (oC)
Magnesium Flouride MgF2 Ionic 1261 2512
Sodium Chloride NaCl Ionic 801 1686
Calcium Iodide CaI2 Ionic 784 1373
Iodine MonoChloride ICl Covalent 27 370
Carbon tetrachloride CCl4 Covalent -23 350
Hydrogen Flouride HF Covalent -83 293
Hydrogen Sulfide H2S Covalent -86 212
Methane CH4 Covalent -182 109
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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

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Covalent Bonds

• Sharing of electrons between 2 nonmetals

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Non-Polar Covalent

• Electrons are shared equally
• Equal distribution of electrons
• All diatomic molecules have non-polar covalent
  bonds

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Nonpolar Covalent Examples

• Flourine (F2)
• Hydrogen (H2)

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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
• The greater the electronegativity difference, the
  more polar the bond is

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Polar Covalent Examples

• HCl
• H2O

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Dipoles

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

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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-
33
Dipole Examples

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

34
Properties of Molecular Substances (Covalent
Compounds)

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

35
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

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Prefixes
Mono- 1 Hexa- 6
Di- 2 Hepta- 7
Tri- 3 Octa- 8
Tetra- 4 Nona- 9
Penta- 5 Deca- 10

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

37
Examples

1. CCl4
2. H2O
3. NO
4. N2O5
5. BBr3

38
Structural Formulas

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

39
Single Bonds

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

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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

43
Network Solids

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

44
Network Solids

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

45
Metallic Bonding

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

46
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)
• All solids at STP (except Hg)
• Malleable, Ductile

47
VSEPR Theory

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

48
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

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Tetrahedral
50
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
51
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

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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Bent
54
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

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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
• Polar molecules will have stronger IMFs than
  nonpolar molecules
• The greater the polarity the stronger the IMF

56
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

57
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

58
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

59
Hydrogen Bonds

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

60
H-bonds Examples

• 1. Which sample has Hydrogen Bonds?
• a. H2 b. HF c. F2 d. HCl
• 2. Rank in order from strongest (1) to weakest
  (3).
• a. Hydrogen Bonds
• b. Covalent Bonds
• c. Dipole-Dipole Attractions

61
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)

62
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

63
Van Deer Waals Forces

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

64
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