Unit 4 - MoLECULES - PowerPoint PPT Presentation

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Unit 4 - MoLECULES

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Title: Unit 4 - MoLECULES


1
Chemical Bonds Continued
  • Unit 4 - MoLECULES

2
NaCl vs. CO2
  • What do you already know?
  • Imagine!
  • Close your eyes and picture a piece of salt. Now
    zoom in..what does it look like at the atomic
    level?
  • Now imagine carbon dioxide what does it look
    like at the atomic level?
  • How are these two compounds different?

3
How do non-metals bond with each other?
  • Recall non-metals have______(high or low)
    ionization energy when compared to metals.
  • Reason
  • Result
  • What is electronegativity? Why does it exhibit a
    distinct periodic trend?
  • How do nonmetals bond?

4
Electronegativity - Linus Pauling
5
Covalent bonding
  • Atoms are bonded because both nuclei () are
    attracted to the same electrons (-).

6
Covalent vs Ionic
7
Using electronegativities
  • Electronegativity Difference /Bond Character
  • gt 1.7 and above / ionic
  • 0.4 - 1.7 / polar covalent
  • 0 -0.4 /nonpolar covalent

8
What is a molecule?
  • Two or more atoms covalently bonded to make a
    neutral particle is a unit called a molecule.
  • A polyatomic ion (i.e. NO3-) is very similar to a
    molecule, except that it has a charge.
  • Covalent compounds AND most non-metal elements
    are composed of molecules.
  • H2O CO2 O2 P4

9
MORE KEY TERMS
  • All diatomic molecules are NONPOLAR COVALENT
  • Practice lewis dot structure
  • Electronegativity tendency of an atom to
    attract shared electrons to itself! (greatest
    attraction for electrons)

10
What is a Lewis symbol?
  • Lewis symbols
  • are simple pictures of atoms
  • are used to represents covalent bonds

Lewis Symbol
Ne
Each dot is a valence electron
Gilbert N. Lewis
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12
Molecular Representations
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14
Diatomic nonpolar molecules
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16
Polar covalent vs. nonpolar covalent
17
What determines the structure of molecule?
  • Octet rule - atoms usually form covalent bonds
    with each other resulting in a total of 8 valence
    electrons around each atom.
  • What is special about eight e-?
  • Why usually but not always 8?
  • Whats the actual REASON that non-metals share
    electrons?

18
Drawing Lewis Dot Structure
  • 1. Write symbols of the elements.
  • - if 3 or more elements figure out which one is
    the central one ( typically C , N, P, S
    sometimes O) If all of these are present
    usually C is the central one.
  • 2. Determine total of valence electrons
  • Dont forget about adding or subtracting an
    electron/s if you are dealing with the polyatomic
    ions
  • 3. Use a single bond to connect each atom
    together
  • -gt then fill in the remaining electrons
    around the atoms to complete the octet rule
  • - if total of electrons wont fulfill octet
    rule double or triple bonds are
    necessary

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Example Lewis structure for water
  • Formula of water is H2O
  • Total valence electrons
  • Lewis structure to obey octet rule
  • Practice on your own PCl3 and SF2

25
Lewis structures have limitations
  • What does a Lewis structure show us about a
    molecule?
  • What does it NOT show?

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CF4
31
CF4
32
NH3
33
NH3
34
What about the ammonium ion? NH4
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N2
38
N2
39
nomenclature for molecules
  • Rules for naming covalent molecules
  • 1. Use prefixes to indicate the of atoms
    present
  • EXCEPTION
  • Never use mono for the first element in the
    molecule
  • 2. Same rules for ide ending for the last
    element in molecule
  • Prefixes
  • Mono -1 hexa -6
  • Di- 2 hepta -7
  • Tri -3 octa -8
  • Tetra- 4 nona -9
  • Penta- 5 deca -10

40
Practice!
  • Name these compounds
  • PCl3
  • H2O
  • N2O4
  • SF6
  • Write formulas for these compounds
  • Diphosphorus pentoxide
  • Carbon tetrachloride

41
VSEPR theory
42
Linear shape (2/0)
43
Trigonal planar 3/0
44
Tetrahedral 4/0
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46
Trigonal Pyramidal(3/1)
47
Bent 2/2 or 2/1
48
Bonds to central atom lone pairs on central atom Molecular geometry Bond Angle example
2 0 Linear 180
2 1 Bent 120
2 2 Bent 104.5
3 0 Trigonal Planar 120
3 1 Trigonal Pyramidal 107
4 0 Tetrahedral 109.5
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Molecular Structures and the Periodic table
51
Summary of 4 major structural units of C.
52
VSEPR video
  • http//www.youtube.com/watch?vi3FCHVlSZc4NR1

53
Intramolecular Forcesvs. Intermolecular
  • Intermolecular Forces
  • Hydrogen bonding
  • Dipole forces
  • London forces (dispersion forces)

54
Intra vs Inter
  • Intramolecular forces (within a compound)
  • Covalent and ionic bonds
  • Intermolecular forces (between compounds)
  • IMFs Intermolecular Forces
  • Dipole Forces
  • Hydrogen Bonding
  • LDFs London Dispersion Forces

55
Bond Polarity vs. Molecule Polarity
  • Bond Polarity results from unequal sharing of
    the electrons in the covalent bond. Use the
    electronegativity differences to figure out how
    polar the bond is.
  • Molecule Polarity a molecule is polar if it has
    1 or more polar bonds and its shape does not
    cancel out the polarity.
  • Nonpolar covalent molecules (nonpolar)
  • Ex All diatomics such as N2, H2, etc
  • Polar covalent molecules (polar)
  • Ex water, carbon monoxide, etc

56
Intramolecular ForceTutorial Video Advanced,
But Slow and Methodical
57
Using electronegativities
  • Electronegativity Difference /Bond Character
  • gt 1.7 and above / ionic
  • 0.41 - 1.7 / polar covalent
  • 0 -0.4 /nonpolar covalent
  • Dipole molecule -

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Polar Molecule notations/representations
60
Intermolecular Forces Hydrogen bonding
61
Hydrogen bonding
62
Hydrogen bonding characteristics
  • 1. Must involve a hydrogen that is bonded to
    highly electronegative element (O, F, N)
  • 2. The slight positive on the hydrogen is
    attracted to a neighboring molecules nonbonding
    electron pair.

63
Dipole Forces (Vander wall)
  • Dipoles molecules that have separate centers of
    partial negative and partial positive charges.
  • Note dipole forces are only 1 as strong as an
    ionic bond attraction.

64
London dispersion forces (LDFs)
  • They are small, transient , attractive forces
    between NONPOLAR molecules
  • Larger or heavier atoms typically exhibit
    stronger dispersion forces than smaller, lighter
    molecules
  • http//www.youtube.com/watch?v3t1Jn_jrsQk
  • http//dl.clackamas.edu/ch104/lesson9molecular_pol
    arity.html

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Ionic vs molecular compounds
  • Melting point
  • Conductivity

68
Intermolecular forces and boiling points!
  • The GREATER The force of attraction between
    molecules the higher the melting or evaporation
    point will be.
  • Think about Ionic bonding..
  • Which would melt the quickest when heated and
    why?
  • Aluminum nitride
  • Sodium chloride
  • Calcium oxide
  • Now what about covalent compounds..

69
Boiling points of ionic vs. molecular
  • Melting Points and Boiling Points of Substances
    with Similar Formula Weights Substance FW (g/mol)

Covalent molecules b.p. (Celsius) Ionic compounds b.p. (Celsius)
CS2 46.0 NaF 1695
CH4O 64.7 CaCO3 825
propane -42.1 NaCl 801
CO2 -57 MgSO4 1125
ethanol 78.5 TiO 1750
Glucose 146
H2O 100
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