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Through the work of
•
•
•
1.
2.
Einstein, Planck, DeBroglie, Bohr,
Schrodinger, Heisenberg we know
Depending on the experiment, light has particle like
behavior (an example…photoelectric effect)
Particles (the electron) has wave like behavior
(example: the emission spectrum)
Our goals in the first part of this chapter is
To describe light and particles in terms of energy
and wavelength, apply DeBroglie’s relationship and
the photoelectric effect
To describe likely positions of where the electron is
located to understand chemical behavior.
• Light is a form of electromagnetic radiation
• Electromagnetic radiation is energy (radiant) that as it travels
consists of electric and magnetic fields
• Travel in waves that can be described by their
wavelength (lamda –l) and frequency (n- nu)
• In a vacuum, travel at the speed of light, c = 3 x 108 m/s
ln= c
White light is made
up of light waves
of different
wavelengths
Disperses light into spectrum of colors. There
are many lines that blend together producing a
continuous spectrum ( 400-700 nm).
Hydrogen
Neils Bohr
Bohr theorized that the lines
in the hydrogen spectrum
corresponded to certain
quantum of energy emitted
or absorbed when the
electron moved from one
“orbit” to another.
• Proposed electrons travel in successively larger
orbits and when an electron jumps from an outer
orbit to an inner one, it emits light.
• Proposed the planetary model - suggested the
electron orbits the nucleus as planets orbit the sun
Spectroscopy
• Method of studying substance that are exposed to exciting
energy. Can be used to identify substances, since each
element emits a unique collection of lines.
• The presence of spectral lines of specific frequencies
suggests that the energy of the electron in an atom is
restricted to a series of discrete values
Max Planck
• Proposed energy is quantized
• Energy can only be lost or gained in packets
• Each packet is called a quantum
Energy of quantum of electromagnetic radiation is
proportional to frequency, DE = hu
Indium compounds give a blueviolet flame test. The atomic
emission responsible for this
blue-violet color has a
wavelength of 451 nm. Calculate
the energy of a single photon of
this wavelength.
Albert Einstein
• Used Planck’s equation to explain
photoelectric effect
Light consists of quanta of energy that behave like tiny
particles of light(photon). Electrons can absorb energy from
photons, but they follow an "all or nothing" principle. All of
the energy from one photon must be absorbed and used to
liberate one electron.
1/2mv2 = hn - f
1/2mv2: Kinetic energy of electron ejected
hn: Energy of incident radiation
f: Work function, minimum energy needed to
eject electron
• The photoelectric work function for Mg is 5.90 x
10-19J. Calculate the minimum frequency of
light required to eject electrons from Mg.
• Light of wavelength 285 nm shines on a piece of
Mg metal. What is the speed of the ejected
electron?
Albert Einstein
• Used Planck’s equation to explain
photoelectric effect
Light consists of quanta of energy that behave like tiny
particles of light(photon). Electrons can absorb energy from
photons, but they follow an "all or nothing" principle. All of
the energy from one photon must be absorbed and used to
liberate one electron.
•In related development, proposed
special theory of relativity
2
E = mc
•
What were the contributions of
DeBroglie? Matter Waves, l = h / mv
• Heisenberg?
• Schrodinger?
Uncertainty Principle – location and momentum
of particle are complimentary; can’t both be
known simultaneously with precision; can’t
specify precise location of particle if it behaves
like a wave
Developed an equation that describes the
wavelike properties of matter, we use the wave
function to express the probability of finding the
electron in a specific region in space (orbital),
Quantum number used to describe the orbital and
the electron
1.
Probability of finding electron at different points is
calculated
• Some points will have higher probability than others
2. If connect all points of high probability, three
dimensional shapes are formed
3. The most probable place to find the electron will be
some place in that shape
Quantum Numbers
• Set of four numbers that completely specify each
electron
1. Prinicipal quantum number, n, corresponds to the
energy level
2. Second quantum number, l, corresponds to the sublevel
within the energy level. It provides information about the
shape of the electron cloud
3. Third quantum number, m, is the orbital quantum
number and describes the orientation in space of orbital
4. Fourth quantum number, s, is the spin quantum
number and describes rotation of an electron on its
axis. Values are ½ or - ½
Sublevel
s
p
d
f
Orientations (Orbitals)
Region around
nucleus where
electrons are likely
to be found. Each
orbital holds two
electrons that have
opposite spin
6d __ __ __ __ __
5f __ __ __ __ __ __ __
7s __
6p __ __ __
5d __ __ __ __ __
4f __ __ __ __ __ __ __
6s __
5p __ __ __
4d __ __ __ __ __
5s __
4p __ __ __
3d __ __ __ __ __
4s __
3p __ __ __
3s __
2p __ __ __
2s __
1s __
Atomic Orbitals
• Region around
nucleus where
electrons are
likely to be
found
Energy Level
Sublevels
1
s
2
s,p
3
s, p, d
4
s, p, d, f
• Each orbital holds two electrons that
have opposite spin
Follow 3 rules to configure the electrons
1. Aufbau Principle - electrons fill orbitals
starting at the lowest available (possible)
energy states before filling higher states
2. Pauli Exclusion Principle - two electrons
cannot share the same set of quantum
numbers within the same system. Therefore,
there is room for only two electrons in each
orbital and the electrons have opposite spin.
3. Hund’s Rule – in equal energy orbitals,
arrange the electrons to achieve the maximum
number of unpaired electrons.
DO NOW
• Draw the orbital notation for arsenic.
How many unpaired electrons does it
have? Compare paramagnetic and
.
diamagnetic. Paramagnetic has net
unpaired spins and it attracted by a
magnet. Diamagnetic does not.
• Write the electron configuration for
manganese, Br -, Ba2+
• Depict an atom (Bohr Model) of
Uranium-235
• Take out a periodic table and paper only.
1. Write out the electron configuration for
osmium-191
2. Write out the electron configuration for
• A) Na and Li
• B) O and S
• C) He, Ne, Ar
• For each grouping in number 2, do you
notice any similarities?
Write out the electron
configuration
for
• Lithium
• Oxygen
• Cesium
Abbreviated Electron
Configuration
• What is a trend?
• Atomic Radius
•Ionic Radius
• Ionization Energy
•Electron Affinity
Atomic Radius
• Distance from nucleus to outermost
electron
Down a Group-
Increases
Across a Period
Decreases
–
Ionic Radius
• Distance from nucleus to outermost
electron in an ion
Positive ion smaller than atom
+
<
ATOM
Negative ion larger than atom
ATOM
<
-
Ionization Energy
• Minimum energy required to lose an
electron and form a positive ion
Down a Group-
Decreases
Across a Period
Increases
–
Successive Ionization Energies
• Look at the chart on p. 310
• If you scan Magnesium’s ionization
energies a large jump is observed between
the 2nd and 3rd. Explain.
• Where would you predict a large jump in
successive ionization energies for Sr?
Practice!
• 1. Compare the sizes of… (Explain your trend)
a.Chlorine and magnesium
b.Nitrogen and bismuth
2. What is the charge of the ion formed by
a. nitrogen? Also compare the radius of the
nitrogen atom to the nitride ion..
b. magnesium? Compare magnesium ion and
atom radii.
3. Compare ionization energy of____. Which is
larger and why? a. strontium and antimony
b. sodium and francium
Electron Affinity
• Energy to gain an electron and form a
negative ion
Down a Group-
Increases
Across a Period
Decreases
–
Electronegativity
• Ability to attract electrons in a chemical
bond
Down a Group-
Decreases
Across a Period
Increases
–