Covalent Bonding and Naming

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Covalent Bonding and Naming
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I. Types of Covalent Bonds

• l. Nonpolar covalent bond-a covalent bond in
  which the bonding electrons are shared equally
• 2. Polar covalent bond-a bond is which the bonded
  atoms do NOT share the bonding electrons equally.
  

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Types of Covalent Bonds

• If the difference in electronegativity between
  two bonded atoms is from 0.3 to 1.7, a polar bond
  will exist
• If the difference in EN is less than 0.3 then the
  bond is nonpolar covalent.

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Practice

• H and Cl
• F and Br
• S and I
• O and H

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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 energy
  level.

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Multiple Bonding

• single bond -two atoms share one pair of
  electrons
• double bond -two atoms share two pair of
  electrons
• triple bond -two atoms share three pair of
  electrons

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Drawing Lewis Dot Structures

• Add up the TOTAL number of valence electrons in
  the substance
• Decide what is the central atom. The central
  atom is the one that is least represented. (or
  the least electronegative)
• Arrange the dots so that each atom has an octet

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Practice

• 1. Br2
• 2. CCl4
• 3. H2O
• 4. NH3
• 5. O2
• 6. N2

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Exceptions to the Octet Rule

• Some atoms have less than an octet. Example
  Hydrogen only needs 2 electrons surrounding it
  and boron only needs 6.
• H2
• BF3

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Exceptions to the Octet Rule

• Some atoms have more than an octet (One reason
  is because of bonding d orbitals as well as s and
  p orbitals.) Example Sulfur can have up to 12
  electrons surrounding it.
• SF6

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Resonance

• a concept in which two or more Lewis structures
  for the same arrangement of atoms (resonance
  structures) are used to describe the bonding in a
  molecule or ion. To show resonance, a
  double-headed arrow is placed between a
  molecules resonance structures. Example O3

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Molecular Shapes

• The Valence Shell Electron Pair Repulsion Theory
  (VSEPR) is used to predict the shape of
  molecules.
• The VSEPR theory states that the bonding and
  nonbonding pairs of electrons will arrange
  themselves so that the repulsive forces between
  them are at a minimum. Molecules will adjust
  their shapes so that the nonbonding electrons are
  as far apart as possible.

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Molecular Shapes

• The number of shared and unshared pairs of
  electrons determines the shape of the molecule.

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Linear

• The bond angle is 180º .
• 2 atoms bonded to the central atom
• 0 lone pairs
• Ex CO2

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Bent

• The bond angle is 90º
• 2 atoms bonded to the central atom
• 2 lone pairs
• Write the Lewis structure for SF2

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Trigonal Planar

• The bond angle is 120
• 3 atoms bonded to the central atom
• 0 lone pairs
• Write the Lewis structure for BCl3

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Trigonal Pyramidal

• The bond angle is 107
• 3 atoms bonded to the central atom
• 1 lone pairs
• Write the Lewis structure for NH3

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Tetrahedral

• The bond angle is 109.5
• 4 atoms bonded to the central atom
• 0 lone pairs
• Write the Lewis structure for CH4

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Trigonal bipyramidal

• The bond angle is 120
• 5 atoms bonded to the central atom
• 0 lone pairs
• Write the Lewis structure for PCl5

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Octahedral

• The bond angle is 120
• 6 atoms bonded to the central atom
• 0 lone pairs
• Write the Lewis structure for SF6

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Polarity

• A. nonpolar covalent bonds EN difference is
  less than 0.3
• There is equal sharing of electrons
• B. polar covalent bonds EN is between 0.3 and
  1.7
• There is unequal sharing of electrons

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Practice

• Determine whether each of the following is a
  polar covalent bond, nonpolar covalent, or ionic
  bond.
• 1. N-O bond 2. Cl-Cl bond
• 3. C-S bond 4. S-O bond
• 5. P-O bond 6. Na-Cl bond

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Determining Polarity

• In a polar molecule there is an uneven
  distribution of charge. One end of the molecule
  is more negative or positive than the other
• In a nonpolar molecule there is an even
  distribution of charge.

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Determining Polarity

• Molecules with nonbonding pairs of electrons on
  the central atom are polar.
• H2O
• NH3

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Determining Polarity

• The polarity of molecules with no nonbonding
  pairs depends on the atoms bound to the central
  atom.
• If the surrounding atoms are identical, the
  molecule is nonpolar. This is because the bond
  dipoles cancel.
• Example CH4
• When the atoms surrounding the central atom are
  different, the molecule is polar.
• Example BBrI2

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Practice

• SO2
• H2S
• CO2

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• Intermolecular Forces of Attraction (IMFs)

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Van der Waals forces

• The attractive forces between molecules are
  collectively referred to as van der Waals forces.
• 1. There are very strong intermolecular forces
  of attraction between polar molecules and weak
  intermolecular forces of attraction between
  nonpolar molecules.
• 2. Many physical properties of a substance, such
  as freezing point and melting point, depend on
  the strength of the intermolecular forces of
  attraction.

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Dipole-Dipole Forces

• 1. Dipole-dipole forces exist between polar
  molecules. The negative dipole of one molecule
  attracts the positive dipole of another molecule.
• Example HCl
• 2. Dipole-dipole forces are weaker than chemical
  bonds.

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Hydrogen Bonding

• 1. Hydrogen bonding is a special type of
  dipole-dipole force. Since no electrons are
  shared or transferred, hydrogen bonding is not a
  chemical bond.
• 2. Hydrogen bonding always involves molecules
  containing hydrogen that is chemically bonded to
  a highly electronegative atom of small atomic
  size, specifically nitrogen, oxygen or fluorine
  Example H2O

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London Dispersion Forces

• 1. London dispersion forces are most noticeable
  in nonpolar molecules
• 2. London dispersion forces arise from the motion
  of electron clouds . From the probability
  distributions of orbitals, it is concluded that
  the electrons are evenly distributed around the
  nucleus. However, at any one instant, the
  electron cloud may become distorted as the
  electrons shift to an unequal distribution. It is
  during this instant that a molecule develops a
  temporary dipole.This temporary dipole introduces
  a similar response in neighboring molecules, thus
  producing a short-lived attraction between
  molecules

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Strength of Forces of Attraction from Weakest to
Strongest

• 1. Chemical Bonds
• a. Nonpolar Covalent
• b. Polar Covalent
• c. Ionic
• 2. Van der Waals Forces
• a. London Dispersion Forces
• b. Dipole-Dipole Forces
• c. Hydrogen Bonding

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Comparing Boiling Points

• 1. When molar masses are similar, substances with
  stronger intermolecular forces of attraction have
  higher boiling and freezing points.
• 2. In order to compare boiling and freezing
  points of a substances with similar molar masses
  you must
• a. Identify the structural formula
• b. Determine the polarity of the molecule.
• c. Determine the dominant type of Van der Waals
  forces
• 3. Which of the following molecules has the
  higher boiling point CH4 or NH3? Explain your
  answer.

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Naming and Writing Covalent Compounds

• Writing Formulas for Binary Molecular Compounds
  those containing 2 nonmetals. Use the prefix
  naming system - know theses prefixes

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Naming and Writing Covalent Compounds

• mono one
• di two
• tri three
• tetra four
• penta - five
• hexa six
• hepta seven
• octa - eight
• nona - nine
• deca ten

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Practice

• nitrogen tetrasulfide
• carbon dioxide
• oxygen monofluoride
• sulfur hexachloride
• trioxygen decanitride
• tetrafluorine monophosphide
• hexafluorine nonasulfide
• heptabromine octanitride

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Naming Binary Molecular Compounds

• The less electronegative element is given first.
• The second element is named by combining a prefix
  indicating the number of atoms contributed by the
  element to the root of the name of the second
  element and then adding ide to the end.

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Practice

• CCl4 _________________________
• NF3 _________________________
• PBr5 _________________________
• SF6 _________________________
• SO3 _________________________
• PCl5 _________________________
• N2O _________________________ PF6
  _________________________

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• Naming Acids

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• Binary or hydrohalic acids
• Name hydro____ic acid
• Ex HF ____________________
• HCl ____________________
• HBr ____________________
• HI ____________________
• H2S ____________________

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• Oxyacids contain a polyatomic ion
• If the polyatomic ion ends ate the acid will
  end in -ic
• HNO3 ____________________
• H3PO4 ____________________
• H2SO4 ____________________
• H2CO3 ____________________
• HC2H3O2 ____________________

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• Oxyacids contain a polyatomic ion
• If the polyatomic ion ends ite the acid will
  end in -ous
• HNO2 ____________________
• H3PO3 ____________________
• H2SO3 ____________________

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Metallic Bonds

• Each metal donates its valence electron(s) to
  form an electron cloud
• This leaves positive particles which are
  "cemented" together with the negative electron
  cloud, often called a sea of electrons.