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CHEMICAL
CHEMICAL
BONDING
BONDING
Chemical Bonding
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How is a molecule or
polyatomic ion held
together?
Why are atoms distributed at
strange angles?
Why are molecules not flat?
Can we predict the structure?
How is structure related to
chemical and physical
properties?
How is all this connected with
the periodic table?
Cocaine
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ATOMIC
STRUCTURE
ELECTROMAGNETIC
RADIATION
Periodic Table & Chemistry
Li,
Li, 3eLi+
Na,
Na, 11eNa+
Mg, 12eMg2+
Al, 13eAl3+
C, 6eCH4
Si,
Si, 14eSi4+, SiH4
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Electromagnetic
Electromagnetic Radiation
Radiation
• Most subatomic particles behave as
PARTICLES and obey the physics of
waves.
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3
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Electromagnetic
Electromagnetic Radiation
Radiation
wavelength
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Electromagnetic
Electromagnetic Radiation
Radiation
Visible light
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Electromagnetic
Electromagnetic Radiation
Radiation
Long wavelength --> small frequency
Short wavelength --> high frequency
• Waves have a frequency
ν
• Use the Greek letter “nu
”, , for frequency,
“ nu”,
and units are “cycles per sec”
Amplitude
λ ν = c
wavelength
• All radiation:
•
where c = velocity of light = 3.00 x 10 8 m/sec
• Long wavelength --> small frequency
• Short wavelength --> high frequency
Node
Ultaviolet radiation
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Electromagnetic
Electromagnetic Radiation
Radiation
Red light has λ = 700 nm.
nm. Calculate the
frequency.
Quantization of Energy
Max Planck (1858-1947)
Solved the “ultraviolet
catastrophe”
1 x 10 -9 m
700 nm •
= 7.00 x 10-7 m
1 nm
Freq =
3.00 x 10 8 m/s
7.00 x 10 -7 m
= 4.29 x 10
14
sec
increasing
frequency
increasing
wavelength
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Quantization
Quantization of
of Energy
Energy
Energy of radiation is proportional to
frequency
E
E == hh •• νν
-1
h = Planck’s constant = 6.6262 x 10 -34 J•s
An object can gain or lose energy by
absorbing or emitting radiant energy in
QUANTA.
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Photoelectric
Photoelectric Effect
Effect
Quantization
Quantization of
of Energy
Energy
E
E == hh •• νν
Light
Light with
with large
large λλ (small
(small νν)) has
has aa small
small E.
E.
Light
) has
Light with
with aa short
short λλ (large
(large νν)
has aa large
large E.
E.
A. Einstein (1879-1955)
• Experiment demonstrates the particle
nature of light. (Figure 7.6)
• Classical theory said that E of ejected
electron should increase with increase
in light intensity—not observed!
• No e- observed until light of a certain
minimum E is used.
• Number of e - ejected depends on light
intensity.
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Energy
Energy of
of Radiation
Radiation
Energy of 1.00 mol of photons of red light.
Understand experimental
observations if light consists of
particles called PHOTONS of
discrete energy.
PROBLEM:
PROBLEM: Calculate
Calculate the
the energy
energy of
of 1.00
1.00 mol
mol
of
photons
of
red
light.
of photons of red light.
λλ == 700.
700.nm
nm
νν == 4.29
4.29 xx 10
101414 sec
sec-1-1
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.
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Photoelectric
Photoelectric Effect
Effect
Atomic Line Emission
Spectra and Niels Bohr
E = h•ν
h•ν
10-34
1014
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Bohr’s greatest contribution
to science was in building a
simple model of the atom. It
was based on an
understanding of the
sec-1)
= (6.63 x 10 J•s)(4.29 x 10 sec )
= 2.85 x 10-19 J per photon
E per mol =
(2.85 x 10-19 J/ph
)(6.02 x 1023 ph/
J/ph)(6.02
ph/mol)
mol)
= 171.6 kJ/mol
kJ/mol
This is in the range of energies that can
break bonds.
SHARP LINE EMISSION
SPECTRA of excited
Niels Bohr
(1885-1962)
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atoms.
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Line Emission Spectra
of Excited Atoms
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Line Emission Spectra
of Excited Atoms
• Excited atoms emit light of only
certain wavelengths
• The wavelengths of emitted light
depend on the element.
High E
Short λ
High ν
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The Electric Pickle
• Excited atoms can emit light.
• Here the solution in a pickle is excited
electrically. The Na+ ions in the pickle
juice give off light characteristic of that
element.
Low E
Long λ
Low ν
Visible lines in H atom spectrum are
called the BALMER series.
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Atomic
Atomic Spectra
Spectra and Bohr
Bohr
One view of atomic structure in early 20th
century was that an electron (e-) traveled
about the nucleus in an orbit.
+
Electron
orbit
1.
Any orbit should be possible
and so is any energy.
2.
But a charged particle
moving in an electric field
should emit energy.
End result should be destruction!
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Atomic
Atomic Spectra
Spectra and Bohr
Bohr
Atomic
Atomic Spectra
Spectra and Bohr
Bohr
Bohr said classical view is wrong.
Need a new theory — now called
QUANTUM or WAVE MECHANICS.
MECHANICS.
e- can only exist in certain discrete orbits
— called stationary states.
e- is restricted to QUANTIZED energy
states.
Energy of quantized state = - C/n2
Energy of state = - C/n2
where n = quantum no. = 1, 2, 3, 4, ....
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• Only orbits where n = integral
no. are permitted.
• Radius of allowed orbitals
= n2 • (0.0529 nm)
nm)
• But note — same eqns.
eqns. come
from modern wave mechanics
approach.
• Results can be used to explain
atomic spectra.
If e-’s are in quantized energy
states, then ∆E of states can
have only certain values. This
explain sharp line spectra.
E = -C (1/2 2)
E = -C (1/1 2)
E = -C (1/22)
E = -C (1/12)
n=2
n=1
Calculate ∆E for e- “falling” from high energy
level (n = 2) to low energy level (n = 1).
n=2
∆E = Efinal - Einitial = -C[(1/1 2) - (1/2)2]
∆E = -(3/4)C
n=1
Note that the process is EXOTHERMIC
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Atomic
Atomic
Spectra
Spectra
and
and Bohr
Bohr
ENERGY
Atomic
Atomic
Spectra
Spectra
and
and Bohr
Bohr
Atomic
Atomic Spectra
Spectra and Bohr
Bohr
ENERGY
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E = -C (1/22)
E = -C (1/12)
n=2
n=1
∆E = -(3/4)C
C has been found from experiment (and is now
called R, the Rydberg constant)
R (= C) = 1312 kJ/mol
kJ/mol or 3.29 x 10 15 cycles/sec
so, E of emitted light
= (3/4)R = 2.47 x 10 15 sec-1
and λ = c/n = 121.6 nm
This is exactly in agreement with experiment!
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Atomic
Atomic Line
Line Spectra
Spectra and
and
Niels
Niels Bohr
Bohr
Niels Bohr
(1885-1962)
Bohr’s theory was a great
accomplishment.
Rec’d Nobel Prize, 1922
Problems with theory —
• theory only successful for H.
• introduced quantum idea
artificially.
• So, we go on to QUANTUM or
WAVE MECHANICS
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