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Unformatted text preview: 1 CHEMICAL
BONDING Chemical Bonding 2 How is a molecule or
polyatomic ion held
Why are atoms distributed at
Why are molecules not flat?
Can we predict the structure?
How is structure related to
chemical and physical
How is all this connected with
the periodic table? Cocaine 4 ATOMIC
RADIATION Periodic Table & Chemistry
Li, 3eLi+ Na, 11eNa+ Mg, 12eMg2+ Al, 13eAl3+
Si, 14eSi4+, SiH4 5 6 Electromagnetic Radiation
• Most subatomic particles behave as
PARTICLES and obey the physics of
waves. Page 1 3 7 Electromagnetic Radiation
wavelength 8 Electromagnetic Radiation
Electromagnetic Radiation Visible light Electromagnetic Radiation
Electromagnetic Radiation 9 Long wavelength --> small frequency
Short wavelength --> high frequency • Waves have a frequency ν • Use the Greek letter “ nu”, , for frequency,
and units are “cycles per sec” Amplitude λν=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 10 Electromagnetic Radiation
Red light has λ = 700 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
wavelength 11 12 Quantization of Energy
Quantization of Energy
Energy of radiation is proportional to
frequency E = h•ν
E = h•ν -1 An object can gain or lose energy by
absorbing or emitting radiant energy in
QUANTA. Page 2 h = Planck’s constant = 6.6262 x 10 -34 J•s 13 Quantization of Energy
Quantization of Energy E = h•ν
E = h•ν
Light with large λ ((small ν)) has a small E.
Light with large λ small ν has a small E.
Light with a short λ ((large ν)) has a large E.
Light with a short λ large ν has a large E.
Light Energy of Radiation
Energy of Radiation
Energy of 1.00 mol of photons of red light. Photoelectric Effect
Photoelectric Effect 14 PROBLEM: Calculate the energy of 1.00 mol
PROBLEM: Calculate the energy of 1.00 mol
off photons of red light.
o photons of red light.
λ = 700. nm
λ = 700. nm
ν = 4.29 x 1014 ssec-1
ν = 4.29 x 1014 ec-1 17 . Atomic Line Emission
Spectra and Niels Bohr E = h•ν
10-34 1014 15 Understand experimental
observations if light consists of
particles called PHOTONS of
discrete energy. 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. 16 Photoelectric Effect
Photoelectric Effect 18 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/mol)
= 171.6 kJ/mol
This is in the range of energies that can
break bonds. Niels Bohr
(1885-1962) Page 3 SHARP LINE EMISSION
SPECTRA of excited
atoms. 19 Line Emission Spectra
of Excited Atoms 20 Line Emission Spectra
of Excited Atoms • Excited atoms emit light of only
• The wavelengths of emitted light
depend on the element. High E
High 21 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
Low Visible lines in H atom spectrum are
called the BALMER series. 22 Atomic Spectra and Bohr
One view of atomic structure in early 20th
century was that an electron (e-) traveled
about the nucleus in an orbit. + Electron
Any orbit should be possible
and so is any energy.
But a charged particle
moving in an electric field
should emit energy.
End result should be destruction! 23 24 Atomic Spectra and Bohr
Bohr Atomic Spectra and Bohr
Bohr Bohr said classical view is wrong.
Need a new theory — now called
QUANTUM or WAVE 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, .... Page 4 • Only orbits where n = integral
no. are permitted.
• Radius of allowed orbitals
= n2 • (0.0529 nm)
• But note — same eqns. come
from modern wave mechanics
• Results can be used to explain
atomic spectra. If e-’s are in quantized energy
states, then ΔE 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 ΔE for e- “falling” from high energy
level (n = 2) to low energy level (n = 1). n=2 ΔE = Efinal - Einitial = -C[(1/1 2) - (1/2)2] ΔE = -(3/4)C n=1 Note that the process is EXOTHERMIC
EXOTHERMIC 26 Atomic
and Bohr ENERGY Atomic
and Bohr Atomic Spectra and Bohr
Bohr ENERGY 25 E = -C (1/22) E = -C (1/12) n=2 n=1 ΔE = -(3/4)C
C has been found from experiment (and is now
called R, the Rydberg constant)
R (= C) = 1312 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! 28 Atomic Line Spectra and
Atomic Line Spectra and
Niels Bohr Niels Bohr
(1885-1962) Bohr’s theory was a great
Rec’d Nobel Prize, 1922
Problems with theory —
• theory only successful for H.
• introduced quantum idea
• So, we go on to QUANTUM or
WAVE MECHANICS Page 5 27 ...
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