Molecular Nomenclature and Geometry

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Molecular Nomenclature and Geometry

• Chemistry Text Ch 6.1,6.2,6.5,16.1-16.3

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Covalent Bonding

• Covalent bonding entails a sharing of electrons.
• Covalent bonding usually occurs between
  nonmetals.
• Form individual molecules.
• Shapes of molecules determines properties.

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Properties of Covalent Molecules

• Gases, liquids, or solids (made of molecules)
• Low melting and boiling points
• Poor electrical conductors in all phases
• Many soluble in nonpolar liquids but not in water

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Covalent Bonds
attain the octet or full valence by sharing
pairs of valence electrons.
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Covalent Bonds

• For every pair of electrons shared between two
  atoms, a single covalent bond is formed. 
• Some atoms can share multiple pairs of electrons,
  forming multiple covalent bonds. 

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Building a Dot Structure

• Hydrogen - H2
• 1. How many valence electrons? (1)
• 2. Add up the number of valence electrons that
  can be used.
• H 1 and H 1
• Total (1 1) 2
• 2 electrons
• 3. Put dots together
• With dots- HH

Electrons
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Building a Dot Structure

• Hydrogen - H2
• With dots- HH
• 4. When two dots are together,
• replace with a line.
• HH H-H

Electrons
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Building a Dot Structure

• Ammonia, NH3
• 1. Number of Valence electrons
• H 1 N5
• Decide on the central atom never H.
• 2. Add up the number of valence electrons that
  can be used.
• H 1 and N 5
• Total (3 x 1) 5
• 8 electrons / 4 pairs

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Building a Dot Structure

• 3. Form a single bond between the central atom
  and each surrounding atom (each bond takes 2
  electrons!)

..
4. Remaining electrons form LONE PAIRS to
complete the octet as needed (or duet in the case
of H).
3 BOND PAIRS and 1 LONE PAIR.
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Building a Dot Structure
..

• 5. Check to make sure there are 8 electrons
  around each atom except H. H should only have 2
  electrons. This includes SHARED pairs.

6. Also, check the number of electrons in your
drawing with the number of electrons from step 2.
If you have more electrons in the drawing than
in step 2, you must make double or triple bonds.
If you have less electrons in the drawing than in
step 2, you made a mistake!
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Simplest Organic molecule
? The Hs can be replaced with any group 17
Halogen(s) to make a multitude of molecular
compounds
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• Diatomic Elements - There are seven elements that
  exist as diatomic molecules in which two atoms of
  the same element bond together.
• They are bromine, iodine, nitrogen, chlorine,
  hydrogen, oxygen, and fluorine. If the symbols
  are written for these elements in the order
  given, they spell out Br I N Cl H O F.
• Whenever these elements appear as free elements
  (by themselves) in a chemical equation, they MUST
  have a subscript "2"
• Ex. Br2 I2 N2 Cl2 H2 O2 F2

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Double Covalent Bond

• Oxygen (which has six valence electrons) needs
  two electrons to complete its valence shell. Two
  oxygen atoms form the compound O2, they share two
  pairs of electrons, forming two covalent bonds
  (double bond).  

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Triple Covalent Bonds

• Nitrogen (which has five valence electrons) needs
  three electrons to complete its valence shell.
  Two nitrogen atoms form the compound N2, they
  share three pairs of electrons, forming three
  covalent bonds (triple bond). 

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• Covalent Bonds
• Draw the Lewis structures for each atom, draw
  circles to show the electrons that are shared,
  and then write the bond structure and chemical
  formula.
• Fluorine Fluorine
• (B) 3 Hydrogen 1 Phosphorus
• (C) 2 Hydrogen 1 Sulfur
• (D) 1 Carbon 2 Oxygen

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• lowest energy state- allows four bonds by
  electron promotion

• Excited state

Hybridization
4 valence e-/sp3
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VSEPR

• VSEPR stands for Valence Shell Electron Pair
  Repulsion.  
• Basically, the idea is that covalent bonds and
  lone pair electrons like to stay as far apart
  from each other as possible under all
  conditions. 
• This is because covalent bonds consist of
  electrons, and electrons don't like to hang
  around next to each other much because they have
  the same charge (like charges repel).

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VSEPR explains why molecules have their shapes. 

• If carbon has four atoms stuck to it (as in CH4),
  these four atoms want to get as far away from
  each other as they can.  This isn't because the
  atoms necessarily hate each other, it's because
  the electrons in the bonds hate each other. 
  That's the idea behind VSEPR.

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Methane is Tetrahedral
Sp3 hybridized carbon 4 equivalent C-H bonds
(s-bonds) All purely single bonds are called
s-bonds
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Molecular Geometry
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Polarity

• Depending on the percent covalent vs. ionic
  characteristic of the bond, molecular compounds
  can have polar covalent or nonpolar covalent
  bonds
• The higher the percentage of ionic
  characteristic the more polar the bond will be

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Naming Covalent Compounds
Covalent compounds are named by adding prefixes
to the element names.
A prefix is added to the name of the first
element in the formula if more than one atom of
it is present. (The less electronegative element
is typically written first.)
A prefix is always added to the name of the
second element in the formula. The second element
will use the form of its name ending in ide.
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Naming Covalent Compounds
Prefixes
Subscript Prefix
1 mono-
2 di-
3 tri-
4 tetra-
5 penta-
Subscript Prefix
6 hexa-
7 hepta-
8 octa-
9 nona-
10 deca-
Note When a prefix ending in o or a is added
to oxide, the final vowel in the prefix is
dropped.
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Naming Binary Covalent Compounds Examples
1 mono
2 di
3 tri
4 tetra
5 penta
6 hexa
7 hepta
8 octa
9 nona
10 deca
Second element in ide from Second element in ide from
Drop a -o before oxide Drop a -o before oxide
N2S4
dinitrogen tetrasulfide
NI3
nitrogen triodide
XeF6
xenon hexafluoride
CCl4
carbon tetrachloride
P2O5
diphosphorus pentoxide
SO3
sulfur trioxide
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Writing Formulas for Covalent Compounds
The names of covalent compounds contain prefixes
that indicate the number of atoms of each element
present.
If no prefix is present on the name of the first
element, there is only one atom of that element
in the formula (its subscript will be 1).
A prefix will always be present on the name of
the second element. The second element will use
the form of its name ending in ide
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Writing Formulas for Binary Covalent Compounds
Examples
1 mono
2 di
3 tri
4 tetra
5 penta
6 hexa
7 hepta
8 octa
9 nona
10 deca
Second element in ide from Second element in ide from
Drop a -o before oxide Drop a -o before oxide
nitrogen dioxide
NO2
diphosphorus pentoxide
P2O5
xenon tetrafluoride
XeF4
sulfur hexafluoride
SF6
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Molecular Solid

• Within molecules, covalent bonding holds molecule
  together
• Between molecules, intermolecular forces hold
  molecules together to make them liquids or
  solids.
• Molecules are held in place by intermolecular
  forces
• Properties
• Low to moderate melting point and boiling point
• Soft
• Non conductive

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Intermolecular Forces Intermolecular Forces are
electrostatic forces of attraction that exist
between an area of negative charge on one
molecule and an area of positive charge on a
second molecule.
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• TYPES OF INTERMOLECULAR FORCES
• (only attractions not bonds)
• Hydrogen Bonding ( generally the STRONGEST)
• Van Der Waals
• a. Dipole-dipole (polar molecules)
• b. London dispersion forces (nonpolar
• molecules

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• Two of the intermolecular forces are associated
  with POLAR structures.
• Hydrogen Bonding
• Dipole-dipole Forces
• One of the intermolecular forces is associated
  with NONPOLAR structures.
• London Dispersion Forces

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Hydrogen Bonding These occur between polar
covalent molecules that possess a hydrogen bonded
to an extremely electronegative element,
specifically - N, O, and F.
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Hydrogen Bonding
d -
d

• Hydrogens single electron is pulled toward the
  very electronegative oxygen resulting in
• Large partial charges
• The unshielded nucleus of hydrogen attraction to
  the unshared electron pairs

O
H
d
H
d -
d
O
Hydrogen bond
H
Covalent bond
d
H
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Hydrogen Bonding
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Properities due to Hydrogen Bonding

• Higher than expected melting/boiling points
• More viscous substances (liquids are thicker to
  pour)
• Surface tension an inward pull that minimizes
  the surface area of a liquid.
• Capillary Action

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Hydrogen Bonding
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Hydrogen Bonding Surface Tension
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Only polar covalent molecules have the ability
to form dipole-dipole attractions between
molecules. Polar covalent molecules act as little
magnets, they have positive ends and negative
ends which attract each other.
Dipole-Dipole
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Dipole-Dipole
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London Dispersion Force

• Because of the constant motion of the electrons,
  an atom or molecule can develop a temporary
  (instantaneous) dipole when its electrons are
  distributed asymmetrically about the nucleus.
• The attractive forces are responsible for Bromine
  being a liquid and Iodine a solid at room
  temperature.

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London Dispersion Force

• The electron asymmetry about the nucleus induces
  a temporary attraction between the non-polar
  molecules causing the London Dispersion Force.