Atomic Structure Electron Configurations Periodicity

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					Atomic Structure
Electron Configurations & Periodicity

     Chemistry 40S
     Unit 1
     M. Patenaude
     GPHS Science Dept



                                        1
Introduction
   Atomic structure explains chemical properties
    and patterns of chemical reactivity.
   Chemical reactions involve electrons. Knowing
    where the electrons are, how many, and what
    their energy levels helps explain many chemical
    phenomena.
   Spectroscopy is used to explore atomic
    structure. Because of this, we start with a
    discussion of the nature of light.


                                              Slide 2
The Nature of Light
Particles or Waves??

   Light must be made of
    particles because it…
       travels in a vacuum
       reflects off of objects
       exerts force (on the
        tails of comets)
   Light must consist of
    waves because it…
       reflects like waves
                                  By the end of the 19th century,
       refracts and diffracts    scientists had concluded that
       exhibits interference     light is composed of WAVES!

                                                         Slide 3
Electromagnetic Radiation
   Visible light is just one form of electromagnetic
    radiation
   Light propagates in space as a wave
   In vacuum, speed of light is constant and given
    the symbol “c”, c = 3.00 x 108 m/s
   Light waves have amplitude, frequency, and
    wavelength.
     = wavelength - distance between consecutive
    crests,  = frequency, # of crests that pass a
    given point in one second (SI Unit is s-1 or Hz)
    c = and  = c/
                                              Slide 4
   James Maxwell proposed in the mid-1800’s that light
is composed of perpendicular electric and magnetic waves
                                                   Slide 5
are distances – S.I. unit is the meter




               Which of the above waves has
                  the higher frequency?

                                         Slide 6
The 7 Regions of the E.M. Spectrum




                                Slide 7
Slide 8
Sample Calculation –
wavelength/frequency conversion

   Calculate the frequency of visible light having a
    wavelength of 485 nm?

    Remember to use S.I. units in your calculations!

     =c/
        = (3.00 x 108 m/s) ÷ (485 x 10-9 m)
        = 6.55 x 1014 s-1

   What colour of visible light is this?

                                                       Slide 9
Sample Calculation –
wavelength & frequency conversion
   CO2 absorbs light with a wavelength of 0.018 mm.
    Which frequency is this?
          c 3.00  108 ms1
                          1.7  1013 s1
           0.018  103 m
          Remember that 1 Hertz (Hz)  1 s-1

   What is the wavelength of WFAE at 90.7 MHz?
          c 3.00  108 ms1
                   6 1
                             3.31 m
            90.7  10 s
                                               Slide 10
Particle Theory of Light
   In 1900, Max Planck turned the world of physics on
    its head by resurrecting the particle theory of light.
   Planck was trying to explain the “ultraviolet
    catastrophe” while studying blackbody radiation.
   He proposed that light is composed of particles
    (quanta) each carrying a fixed amount of energy.
   The amount of energy per quantum is directly
    proportional to the frequency of the light:
                      hc                        34
    E  h    or   E        , h  6.626  10         J s
                       

                                                      Slide 11
Planck’s Equation – sample problem
   The wavelength of maximum visual acuity in
    humans is 550 nm. (green light)
       What is the energy of a single photon having this
        wavelength?

           Ans. 3.61 x 10-19 J per photon

       What is the energy of a mole of photons at 550
        nm?

           Ans. 218 kJ/mol


                                                   Slide 12
Energy, wavelength, frequency

    a)   X-Rays
    b)   Red light
    c)   Green light
    d)   Radio waves


       Rank the above in order of …
        increasing wavelength
        decreasing energy
        increasing frequency

                                       Slide 13
    Atomic Emission Spectra
    “Line Spectra” vs “Continuous Spectra”
    Historic work (1800’s) involving light emitted by pure
     elements in gas phase, subjected to very high voltages.
    The light was passed through a prism or diffraction
     grating to produce a spectrum.
    Each element emitted only certain wavelengths of light –
     a LINE SPECTRUM instead of the more familiar
     “continuous” spectrum seen in the rainbow.
    Balmer and Rydberg described the lines
     mathematically, for hydrogen.
    Hydrogen’s emission spectrum has several lines in the
     visible region – called the “Balmer series”

                                                     Slide 14
A Continuous Spectrum
contains all wavelengths of light – a rainbow




                                                Slide 15
Slide 16
Bohr’s Explanation of Line Spectra

   Neils Bohr developed a mathematical model that could
    explain the observation of atomic line spectra.
   He proposed that electrons orbit the nucleus in certain
    “allowed” orbits - or energy levels. That is, the electron’s
    energy is QUANTIZED (not continuous).
   Working on a model of single-electron atoms, Bohr derived
    an equation to calculate the energy of an electron in the nth
    orbit of such an atom:

                   18  Z2 
    En  2.18  10  2 
                                        where Z = atomic number

                       n              and n = electron energy level


                                                          Slide 17
Bohr’s Quantum Model for
Hydrogen
                  The electron in hydrogen occupies
                   discrete energy levels.
                  The atom does not radiate energy
                   when the electron is in an energy
                   level.
                  When an electron falls to a lower
                   energy level, a quantum of radiation
                   is released with energy equal to the
                   difference between energy levels.
                  An electron can jump to higher
                   energy levels if the atom absorbs a
                   quantum of radiation with sufficient
                   energy.

                                           Slide 18
The Hydrogen Line Spectrum




                       Slide 19
Bohr’s Model: Calculations

   Show that an              Step 1: Calculate E4 & E2
    electron transition           E4 = -1.36 x 10-19 J
    from n = 4 to n = 2           E2 = - 5.45 x 10-19 J
    results in the
    emission of visible
    light. Calculate the      Step 2: Find DE42
    wavelength of the             DE = E2 – E4 = 4.09 x 10-19 J
    light emitted.
                              Step 3: Calculate 
                                  = 4.86 x 10-7 m = 486 nm


                                                           Slide 20
What’s Wrong with Bohr’s Model??

   Although it works great        Ultimately, the failure of
    for single-electron             Bohr’s model lay in the
    atoms, Bohr’s model             fact that he treated the
    fails for atoms with 2 or       electron as a charged
    more electrons!                 particle orbiting the
   It was a huge leap              nucleus like a planet
    forward, but was                around the sun.
    fundamentally flawed.          Electrons are more
                                    complicated …



                                                       Slide 21
    Wave-Particle Dual Nature of
    Electrons
   Einstein’s famous equation, E = mc2, suggests that energy
    and mass are related – one can convert matter directly into
    energy.
   Louis de Broglie made the leap that if light can behave as
    wave/particles, then so can matter!
   He showed that the wavelength of a baseball is negligible
    (as expected), but that the wavelength of an electron was
    on the order of magnitude equal to that of electromagnetic
    radiation!
   This is what Bohr had missed – an atomic model must
    make use of the wave-nature of electrons to be complete!

                                                      Slide 22
Visualizing the Wave-Nature of
Electrons




                             Slide 23
Heisenberg’s “Uncertainty Principle”
               The more precisely the position is
               determined, the less precisely the
               momentum is known in this instant,
                       and vice versa.


             It is impossible to determine BOTH the
              position and momentum (velocity) of an
              electron - the act of observing the
              electron changes it.
             The electron exists in nature as a wave-
              particle (whatever that is) until an
              experimenter observes it and imposes
              one of the two natures upon it!

                                                Slide 24
"When it comes to atoms,
language can only be used as in
poetry. The poet too, is not nearly
so concerned with describing
facts as with creating images."
       - Niels Bohr to Werner Heisenberg


                                           25
The Heisenburg uncertainty principle is the catch 22 of
physics. Most modern physics and chemistry is largely based on
the study of electrons. In order to view an electron we must use
light to observe it, but in the process of observing we change
the electron’s momentary existence because the light necessary
to view it takes the form of a photon particle. The photon particle
travels at approximately the same velocity as the electron
particle, so that
when it impacts the electron it alters its
position as when two asteroids collide
into each other at similar velocities but
different angles.

Therefore, according to Heisenburg,
it is impossible to know decisively an
electron’s position and momentum.

                                                            Slide 26
Heisenburg proposed that this inability to perceive reality on an atomic
level is not the result of technological ignorance, but rather a fact of the
laws of nature. In other words, Heisenburg believed that no matter how
technologically advanced we become we would never be able to fully
perceive and record atomic reality. This means that an electron’s orbital
path around the nucleus of an atom can only be statistically
approximated.




This is why the quantum atomic model has a cloud of electrons, not a few
electrons following exact elliptical orbits as in Bohr's atomic model. Each
electron in the quantum atomic model is a calculated probability, so the
cloud describes not the exact orbit of millions of electrons, but rather the
probability of the position of only a few electrons at any given time.
                                                                  Slide 27
The Schrödinger Wave Equation
   Schrödinger applied de Broglie’s concept of
    matter waves to the electron
   He explained the quantum energies of the
    electron orbits in the Bohr model of the atom as
    vibration frequencies of electron "matter
    waves" around the atom's nucleus.
   He provided a mathematical equation (wave
    function) to describe electron waves.
   Physicists had difficulty explaining the
    MEANING of the wave function itself, but it
    turns out that its square gives the probability of
    finding the electron in a particular region of
    space in the atom.

                                                         Slide 28
Schrodinger Equation, cont.
   These regions of probability are called orbitals, and this
    concept replaces that of the electron “orbit”.
   Instead of electrons orbiting like planets around the sun,
    Schrodinger’s picture shows an “electron cloud” surrounding
    the nucleus and doesn’t state anything about the path or
    position of electrons within that cloud.
   Orbital “pictures” shown in texts represent regions of space
    where the probability of finding an electron is 90% or greater.




                                                           Slide 29
    Quantum Numbers
   Schrodinger's wave equation used three constants
    (“quantum numbers”) that were used to describe orbitals
    – regions of space where an electron has a probability of
    being found.
   His equation used 3 quantum numbers to describe the
    size & energy, shape, and orientations of the orbitals in
    atoms.
   The three quantum numbers: n, l and ml .
   This set of three numbers provides us with an “orbital
    address” for an electron within an atom – we can
    describe the orbital in which an electron is found if we
    know these three quantum numbers.

                                                     Slide 30
Principal Quantum Number, n
   n can have values > 1, integers only.
   For element atoms on the periodic table in their
    ground states, 1  n  6
   As n increases, energy of electron increases and
    size of orbital increases
   Larger n means greater probability of greater
    distance from the nucleus.
   Electrons with same value of n are said to be in
    the same “electron shell”

                                             Slide 31
Shape Quantum Number, l
   The values for l (lower case letter L, script) are limited by
    the principal energy level – they range from 0 to “n–1”
   Example: If n = 3, l can be 0, 1 or 2


   l gives information about shape of orbital
    l =0        s orbital (found on all principle energy levels)
    l =1        p orbitals (found on level 2 and higher)
    l =2        d orbitals (found on level 3 and higher)
    l =3        f orbitals (found on level 4 and higher)


                                                         Slide 32
Graphs show
 Probability
vs Distance
from nucleus




s Orbitals
               NODE

 on three
  energy
  levels




                      Slide 33
p and d Orbital shapes




                         Slide 34
Slide 35
Orientation Quantum Number, ml
   The values of ml are limited by the value of l for
    a given orbital. They include positive and
    negative integers from -l through +l , inclusive

   For a p orbital, l = 1 , so ml can be: 1, 0, 1
   ml gives information about orientation of
    orbitals in space
   The # of allowed values gives the number of
    orbitals of this type on a given energy level. (3,
    in the case of p orbitals)
                                               Slide 36
 Electron Spin Quantum Number, ms
                                 A fourth quantum number, ms
                                  describes electron spin
                                  (either +½ or –½)
                                 Each electron in atom has a
                                  unique set of these four
                                  quantum numbers.
                                 Electrons in orbitals with same
                                  n and l values are said to be in
A spinning negative charge        the same subshell.
creates a magnetic field.        Electrons with all three numbers
The direction of spin d           the same, n, l , and ml , are in
determines the direction of
                                  the same orbital.
the field.
                                                        Slide 37
Pauli Exclusion Principle
   Electrons have negative charge and repel each
    other. How are the electrons in an atom
    distributed?
   Wolfgang Pauli proposed that no two electrons
    in a given atom can be described by the
    same four quantum numbers!
   The first three quantum numbers determine an
    orbital – therefore spins must be opposite!
   Practical result is that each orbital can hold a
    maximum of two electrons, with opposite spins.

                                            Slide 38
    Concept Check
   Which of the following sets of       What is the maximum
    quantum numbers are NOT               number of electrons in an
    allowed in the hydrogen               atom that can have these
    atoms? For those that are             quantum numbers?
    incorrect, state what is wrong.
                                             n=4
     n       l     ml      ms
                                             n = 5, ml = +1
     2      1       -1     +½
                                             n = 5, ms = + ½
     1      1       0      -½                n = 2, l = 1
     8      7       -6     +½                n = 3, l = 2
     1      0       2      -½                n = 3, ml = 4

                                                              Slide 39
Electron Configurations

   Electrons orbitals are defined by their quantum
    numbers, n, l and ml .
   Each electron in an atom has a unique set of 4
    quantum numbers.
   No two electrons can have the same “address”,
    i.e., the same 4 quantum #’s
   Rules define how multiple electrons will be
    distributed among the possible energy levels


                                              Slide 40
Aufbau Principle
   In the ground state, the electrons occupy the lowest
    available energy levels. An atom is in an excited
    state if one or more electrons are in higher energy
    orbitals.
   In atoms with more than one electron the lower
    energy orbitals get filled by electrons first!
   This is the Aufbau principle, which is named after the
    German word which means “to build up”.
   When one describes the locations of the electrons in
    an atom, start with the lowest energy electron and
    work up to the highest energy electron.


                                                    Slide 41
    Hund’s Rule
   A set of orbitals is said to be “degenerate” if the orbitals
    possess the same energies. For example, all three “2p”
    orbitals on energy level 2 are degenerate. All five “3d”
    orbitals on energy level 3 are also degenerate.
   When filling a set of degenerate orbitals, Hund said that
    electrons should be left unpaired as long as possible so
    as to minimize electron-electron repulsions within the
    orbitals.
   When each degenerate orbital has one electron,
    electrons will then pair, spins opposed, until that sub-
    shell is filled.



                                                       Slide 42
Summary of Distribution Rules
   Electrons distribute to lower energy levels
    until they are filled, before occupying higher
    levels. (Aufbau Principle)
   Electrons will spread out as much as possible
    within a sub-shell. (Hund’s Rule)
   Electrons will pair up, two to an orbital, spins
    opposed, until that sub-shell is filled. (Pauli
    Exclusion)
   Every electron will have unique set of 4
    quantum numbers.

                                               Slide 43
Electronic Configurations
   An electron configuration is a way of describing the
    locations of electrons within an atom.
   Usually written for the ground state of an atom: when its
    electrons occupy orbitals giving the lowest possible
    overall energy for the atom
   Each subshell is designated by “n” and the type of
    orbital. Such as: 1s 3p 4d
   The number of electrons in an orbital is shown as a
    superscript: 1s2 3p3          4d9
   How does one determine the order of filling of the
    orbitals in an atom?? Which orbitals get filled first??


                                                    Slide 44
Orbital Energies in Multielectron
Atoms




                                    Slide 45
The Order of Filling of Orbitals
Simply follow the order of elements in the periodic table!




                                                      Slide 46
Another useful device…
Start at the top and
follow the arrows:

1s,
2s,
2p, 3s,
3p, 4s,
3d, 4p, 5s…




                         Slide 47
Examples of Electron Configurations

   Helium has 2 electrons in the 1s orbital, He: 1s2
   Carbon has 6 electrons, C: 1s2 2s2 2p2
   Calcium has 20 electrons, Ca: 1s2 2s2 2p6 3s2 3p6 4s2
   Calcium cation, Ca2+: 1s2 2s2 2p6 3s2 3p6
   Sulfide anion, S2–: 1s2 2s2 2p6 3s2 3p6

   Notice that the superscripts add up to the total number of
    electrons on the atom or ion! Each of these examples
    represents the “ground state” of the atom/ion because
    the electrons are in the lowest possible energy levels.

                                                     Slide 48
Orbital Diagrams




                   Slide 49
Noble Gas “shorthand” Notation
   The core electrons can be indicated by a short
    form: use the last noble gas to indicate filled
    shells.
   Find the last noble gas before the element under
    consideration. Start the electron configuration
    with the symbol for this noble gas element and
    just tack on the extra electrons:
         Ca: 1s2 2s2 2p6 3s2 3p6 4s2
         Ca: [Ar] 4s2

                                        Ar: 1s2 2s2 2p6 3s2 3p6
                                                      Slide 50
There Have to be EXCEPTIONS
   As atoms get larger, their      Cr: [Ar] 4s1 3d5
    outermost electrons             Cu: [Ar] 4s1 3d10
    become closer in energy
   Large atoms have                    In both cases, one of the 4s
    “exceptions” to the                  electrons is promoted to the
    Aufbau Principle in the              “3d” subshell.
    locations of their valence
    electrons. We will
                                    Fe2+: [Ar] 3d6
    memorize two: Cr and Cu
                                    Cu2+: [Ar] 3d9
   Transition metal cations
    are also strange in their
    electron configurations             Notice that the 4s electrons
                                         are lost before the 3d
                                         electrons!

                                                          Slide 51
Paramagnetism & Diamagnetism

   An atom is said to be            An atom is diamagnetic if
    paramagnetic if it                all its electrons are paired –
    possesses unpaired                it will NOT possess a net
    electrons.                        magnetic field because
   These electrons will result       each electron’s field will be
    in the atom possessing a          cancelled
    net magnetic field               These atoms are NOT
   It will then be attracted         attracted into an external
    into an external magnetic         magnetic field
    field
                                     What is “ferromagnetism”?

                                                         Slide 52
Testing for Paramagnetism
See Kotz & Treichel page 290




                               Slide 53
Electron Configuration
Causes Periodic Trends
   “Periodicity” refers to similarities in behavior and
    reactivity due to similar outer shell electron
    configurations.
   All the Alkali Metals have one unpaired valence
    electron; all the Noble Gases have completely
    filled subshells
   We will examine periodic trends in atomic radius,
    ionization energy, electronegativity, and electron
    affinity

                                                Slide 54
“Valence Electrons”
   Valence electrons are those electrons in the
    highest principal energy level within an atom - the s
    and p electrons in the outermost shell of an atom in
    its ground state.
   These are the electrons involved in forming bonds
    with other atoms during chemical reactions.
   Most common oxidation states (ion charge) for the
    element can be derived from valence electrons
   Completely filled, half filled and empty sub-shells
    have special stability (not sure why!?)


                                                 Slide 55
Valence Configurations
   The “valence configuration” is that part of an
    electron configuration that describes the valence
    electrons.
   For example, sodium’s valence configuration is
    just “3s1”.
   The valence configuration of bromine is “4s24d5”.
    We leave out the “3d10” electrons because they
    are not in the outermost principal energy level.




                                                Slide 56
Periodic Trends: Atomic Size
   Properties of elements and their ions are
    determined by electron configurations
   Individual atoms cannot be accurately
    measured.
   Atomic sizes are determined from bond
    lengths of compounds with various atoms

             Br-Br               Br-C
             2.28 A              1.91 A
             r = 1.14 A          rcarbon = 1.91-1.14
                                            = 0.77 A

                                               Slide 57
Atomic Size, cont.
   Atomic radii increase within a group (column) as
    the principal quantum number of outermost shell
    increases
   Electrons of outermost shells have greater
    probability of being farther from the nucleus
   Atomic size decreases across a row (period)
    from Left to Right, because the effective nuclear
    charge increases.
       There are more and more protons in the nuclei
        attracting more and more electrons!

                                                Slide 58
Atomic Radii:

Atoms get larger
as you move
down a group

Atoms get
smaller as you
 go across a
period




                   Slide 59
Slide 60
Ionization Energies
   First Ionization energy: energy required to
    remove an electron from the ground state
    atom in the gaseous state to make a cation
   1st ionization of Mg:
    Mg  Mg+         I.E. = 738 J/mol
                     [Ne]3s2  [Ne]3s1

   2nd ionization of Mg
    Mg +  Mg2+          1450 J/mol
                         [Ne]3s1  [Ne]
                                            Slide 61
Trend in First Ionization Energies

   First ionization energies decrease down a group
    as radii increase:     Na > K > Cs
   First ionization energies increase across a row
    as radii decrease: Li < F < Ne
   For similar reasons that explained trends in
    atomic radii.
   For a given element, IE2 always > IE1 because of
    the same # protons in nucleus holding onto fewer
    electrons.

                                            Slide 62
Slide 63
Trends in First
Ionization Energies




             Slide 64
First Ionization Energy (cont’d)
   Lowest IE1 are in the             Oxygen has slightly lower
    lower left corner of               IE1 than nitrogen because
    periodic table (Fr and Cs)         it takes more energy to
   Irregularities in IE1 trends       disrupt N’s half-filled
    arise from special                 orbital configuration
    electron configurations.          Boron has slightly lower
   Half filled sub-shells are         IE1 than Beryllium
    especially stable: every           because its first electron
    orbital has only one               is removed from the p-
    electron.                          subshell further from the
                                       nucleus than the first
                                       electron removed in Be

                                                        Slide 65
Electron Affinity
   Energy change when a gas phase atom gains an
    electron to become an anion.
    Cl (g) + e-  Cl– (g) EA = – 349 kJ/mol

   The more negative an EA, the greater the affinity for
    electrons!
   Electron affinities “tend” to become more negative
    across a period and less negative moving down a group
   EA least negative for alkaline earth and noble gas
    elements: they have full or half full s- or p-orbitals and it
    would take a lot of energy to add another electron


                                                        Slide 66
Slide 67
Slide 68
Slide 69
Ion Sizes
   Cations: radius is always smaller than the
    neutral atom.
   Loss of electrons (to make cation) leaves fewer
    electrons attracted to same # of protons
   OR….. loss of outermost shell altogether as in
    group 1A [Ne]3s1  [Ne] for sodium
   Anions: radius is always larger than neutral
    atom.
   Same # protons pulling on more electrons;
    also electrons repel each other.

                                            Slide 70
Slide 71
    Isoelectronic Atoms and Ions
   Isoelectronic atoms or ions have same electron
    configurations – the same number of electrons
    in same orbitals
   Na+ , Ne and F- are isoelectronic: 1s2 2s2 2p6
   For isoelectronic atoms and ions, the species
    with the most protons will be smallest
   It has greater nuclear attraction for the same
    number of electrons, so the electron shells will
    be pulled closer to nucleus.


                                              Slide 72
Bonding in Molecules


   Ionic Bonds
   Polar Covalent Bonds
   Non-Polar Covalent Bonds


                              73
Definitions

   ELECTRONEGATIVITY is the attraction an atom
    has for electrons in a chemical bond.
   A POLAR BOND is one between atoms with
    different electronegativities - one atom attracts
    electrons more strongly than the other. As a
    result, one atom acquires a partial positive
    charge and the other a partial negative charge.




                                              Slide 74
More Definitions


   A DIPOLE is a separation of charge. A polar
    bond is an example of a dipole.
   A DIPOLE MOMENT is the magnitude of a
    dipole - it depends on the size of the charges
    involved.
   A POLAR MOLECULE is one that possesses an
    overall dipole moment.




                                           Slide 75
Classifying Chemical Bonds

   IONIC bonds are characterized as…
       bonds between metals and non-metals
       bonds between cations and anions
       bonds involving a transfer of electrons (from a
        metal to a non-metal)

   The best classification of an ionic bond uses the
    concept of electronegativity.


                                                   Slide 76
Classifying Chemical Bonds


   An ionic bond forms between two atoms with a
    large difference in electronegativity - usually
    stated as being greater than about 1.7 Paulings.

       e.g. NaCl: difference is 2.0 Paulings
       e.g. CuO: difference is 1.7 Paulings




                                                Slide 77
Classifying Chemical Bonds


   COVALENT bonds are usually described as…
       bonds between two non-metal atoms
       bonds involving a sharing of electrons


   Using the concept of electronegativity, we can
    see that there are two types of covalent bonds.




                                                 Slide 78
Classifying Chemical Bonds


   POLAR COVALENT BONDS result from
    unequal sharing of electrons by atoms that
    creates a small dipole.
   This type of bond occurs when there is a small
    but significant difference in electronegativity -
    between 0.4 and 1.7 Paulings.
       e.g. N-H: diff is 0.9 Paulings
       e.g. Cu-S: diff is 0.7 Paulings



                                                Slide 79
Classifying Chemical Bonds


   NON-POLAR COVALENT BONDS result from
    an equal sharing of electrons.
   There is no dipole created since the electrons
    are shared equally.
   This bond forms between atoms with similar (or
    the same) electronegativities - differences
    between 0 and 0.4 Paulings.




                                            Slide 80
Exercise

   Classify the following chemical bonds. Show
    your work.
       Li-F
       Br-Br
       Fe-O
       Al-Cl
       C-S
       C-H


                                            Slide 81
Polarity of Molecules

   A diatomic molecule (e.g. HF) will be polar if its
    bond is polar.
   Larger molecules are polar only when the
    following criteria are met:
       they possess at least one polar bond
       their geometries do not result in the dipoles
        cancelling each other out! (e.g. CO2)




                                                    Slide 82