Transcript Electromagnetic Radiation and Energy

Electromagnetic
Radiation and Energy
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Electromagnetic Radiation:
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Energy traveling through space
Three Characteristics of Waves:
1. Wavelength: (symbolized l)
1. Distance between two consecutive peaks or troughs
in a wave
2. Frequency: (symbolized n)
1. How many waves pass a given point per second
3. Speed: (symbolized c)
1. How fast a given peak moves through space
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Electromagnetic
Radiation and Energy
c=λxν
C = speed of light = 3 x 108 m/s
ν = frequency (s-1 or Hz)
λ = wavelength (m)
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Standing Waves
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Tie down a string at both ends  pluck
Has 2 or more nodes
Distance between nodes is λ/2
Only certain wavelengths are allowed (n x λ/2)
 as is found in atomic theory
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Planck
• Scientists tried to explain relationship between intensity and
wavelength for radiation given off by heated objects
– Electromagnetic radiation color depends on temperature
– Wrongly surmised that the shorter the wavelength the greater the
radiation intensity
– Planck solved the riddle
• He came up with term quantized
– Only certain vibrations with specific frequencies allowed
• Planck’s equation
– Vibrating system energy proportional to frequency of vibration
– E = hn; E in joules, h = (Js) = 6.626 x 10-34 Js = Planck’s constant
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• As temperature
increases…
– Maximum energy
released in visible
spectrum goes towards
UV
• “white hot”
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Einstein’s Photon
• Photoelectric effect
– Electron ejection after light strikes metal surface
• Must be the right frequency
• Automatic door openers
• Einstein
– Light has particle-like properties
• Photons
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Einstein + Planck
• What happens to energy
as frequency increases?
• What happens to energy
as wavelength
increases?
hc
E = h =
λ
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Spectra
• Sunlight yields continuous spectrum
• Energized gaseous elements yield certain
wavelengths
– Line emission spectrum
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Rydberg
• Why did gaseous atoms
emit certain
wavelengths?
• Rydberg equation
– N=3, red line
– N=4, green line 1  R  ( 1  1 ); when n>2 & R = 1.0794 107 m -1
l
22 n 2
– N=5, blue line
• Balmer series
– N=6,7,8
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The Bohr Model of the Atom
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Electron energy quantized
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Electron only occupies certain energy
levels or orbitals
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As “n” increases energy becomes less
negative
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Rhc
; J/atom
n2
Increases
Ground state
If electron in excited energy level
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Potential energy of electron in the nth level = En = -
Only certain amts of E may be
absorbed/emitted
If electron in lowest possible energy
level
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If it didn’t, electron would crash into
protons in nucleus
Excited state
One can calculate energy needed to
Rhc
Rhc
raise H electron per atom from ground E = E f - E i = (- 2 )  (  2 )  984kJ/mol
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state (n=1) to first excited state (n=2)
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Bohr’s Model
• Explains emission spectrum of H
– Movement of electrons from one quantized energy level to
a lower one
• Balmer series
– n > 2 to n = 2
• Visible wavelengths
• Lyman series
– n > 1 to n = 1
• UV (invisible)
• Model only good for one electron system
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Others
• De Broglie
– Electron’s properties (mass &
velocity) related to wave
property (l)
• Schrödinger Wave Equation
– “quantum mechanics”
• quantum numbers
• e- has wave-particle duality
h
λ=
mv
• Heisenberg Uncertainty
Principle
– Probability of e- presence
• Orbital pathways
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Quantum numbers
• Used to solve Schrödinger Wave Equation
• n = principle quantum number
– Principle energy level
• Energy shell
– n1
• l (the letter L) = angular momentum quantum number
– Subshells
• Orbitals
– 0,1,2,…n-1 (s,p,d,f)
• ml = magnetic quantum number
– Orientation of orbitals
• -l...0...l (px, py, pz)
• 2l + 1 (how many orbitals within subshell)
• ms = magnetic spin number = 1/2
– Spin direction of electron in orbital
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Atomic orbital
The probability function that defines the distribution of electron
density in space around the atomic nucleus.
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The s-orbital
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The simplest orbital
The only orbital in the s-subshell
Occurs in every principal energy level
“s” stands for “sharp”
The first energy level only houses this orbital
Can house up to 2 electrons
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The p-orbitals
• Start in second principle energy level (n=2)
• There are three p-orbitals in the p-subshell (see below)
– And one s-orbital
• “p” stands for “principle”
• Can house up to 6 electrons
• Has one nodal surface
– Nodal plane = a planar surface in which there’s zero probability of find an
electron
2px
2py
2pz
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The d-orbitals
• Start in third principle energy level (n=3)
• There are five d-orbitals in the d-subshell
– And one s-orbital
– And three p-orbitals
• Can house up to 10 electrons
• “d” stands for “diffuse”
• Has two nodal surfaces
3dyz
3dxz
3dxy
3dx2-y2
3dz2
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The f-orbitals
• Start in fourth principle energy level (n=4)
• There are seven f-orbitals in the f-subshell
– And one s-orbital
– And three p-orbitals
– And five d-orbitals
• Can house up to 14 electrons
• “f” stands for “fundamental”
• Has 3 nodal surfaces
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Electron configuration
• Electron must be identified as to where it
is located
– Hydrogen:
• One electron in first energy level and s-subshell
– Thus, 1s1 (= Aufbau electron configuration)
• 1 states energy level (n)
• s designates subshell
• Superscript 1 tells how many electrons are in the ssubshell
• Can also use orbital box or line diagrams
– Let’s take a look
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Pauli Exclusion Principle
• An atomic orbital holds a maximum of two electrons
• Both electrons must have opposite spins
• ms = +1/2 & -1/2
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Hund’s Rule
• Electron configuration most stable with
electrons in half-filled orbitals before coupling
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Subshell filling order – not
what one expected
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Using the Periodic Table to
advantage
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Short-hand vs. long-hand
Aufbau electron configuration
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F
Al
Ca
Br
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Exercises
• Give me the Aufbau electron configurations
for:
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Y
Te
Hf
Tl
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Sundry matters pertaining
to d-block metals
• Stability is increased when:
– d-subshell is half-filled (d5)
– d-subshell is completely filled (d10)
• Electrons will be taken from the s-subshell to fill the dsubshell
– But there is a limit
• No more than 2 electrons taken from s-subshell
• Given the above, which subshell electrons will d-block metals
lose first when they ionize?
• So what are the correct electron configurations of Cr and Ag?
• Caveat
– Not all metals follow the above; i.e., take from s-subshell and give to dsubshell
• Ni & Pt, for example
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Sundry matters pertaining to
f-block metals
• Stability is increased when:
– f-subshell is half-filled (f7)
– f-subshell is completely filled (f14)
• Electron will be taken from the d-subshell to
fill the f-subshell
– Eu & Yb
– Am & No
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