Ch. 5- Electrons in Atoms
Unit 7 Targets: The Electronic Structure of Atoms
(Chap 5)
•I CAN Utilize appropriate scientific vocabulary to explain scientific concepts.
•I CAN Perform calculations involving the energy, wavelength and frequency of
electromagnetic waves.
•I CAN Perform calculations to determine the de Broglie wavelength of any
object.
•I CAN Compare and contrast Bohr’s solar system model with Schrodinger’s
wave mechanical model.
•I CAN Predict the movement of electrons from the ground state to an excited
state, back to the ground state.
•I CAN Recognize elements based on their emission spectra. (Flame lab).
•I CAN Generate electron configurations for elements: full electron configuration,
noble gas configuration, orbital diagram.
•I CAN Differentiate relationships between orbitals, sublevels and energy levels.
•I CAN Correlate the relationship between electrons in a single orbital or single
sublevel.
 Electromagnetic
radiation- a form of
energy that exhibits wavelike behavior.
• Wavelength () - length of one complete wave
• Frequency () - # of waves that pass a point
during a certain time period
 SI UNIT for frequency is hertz (Hz) = 1 wave / s
• Amplitude (A) - distance from the origin to the
trough or crest

crest
A

greater
amplitude
origin
(intensity)
A
trough
greater
frequency
(color)
• To understand the electronic structure of atoms
we must understand light and how it is emitted or
absorbed by substances.
• We will examine visible light a type of
Electromagnetic Radiation (EM) which carries
(radiant) energy through space (speed of light) and
exhibits wavelike behavior.
• Also need to think of light as particle, to help
understand how EM radiation and atoms interact
H
I
G
H
E
N
E
R
G
Y
L
O
W
E
N
E
R
G
Y
Move through a vacuum at the ‘speed of
light’ 3.00 x 108 m/s
Behaves like waves that move through
water, which are the result of a transfer of
energy to the water (from a stone),
expressed as up and down movement of
water
Both electric and magnetic properties
H
I
G
H
L
O
W
E
N
E
R
G
Y
red
R O Y
G.
orange
green
yellow
B
blue
I
indigo
V
violet
E
N
E
R
G
Y
Wave Speed = (distance between wave peaks) x (frequency)
=
(wavelength)
x (frequency)
EM radiation moves through a vacuum at the “speed of light”
3.00 x 108 m/s also called c.
A lower energy wave (infrared and red) has a longer
wavelength() and lower frequency(f)
A higher energy wave (blue - violet) has a shorter
wavelength() and higher frequency(f).
 Frequency
& wavelength are inversely
proportional
c = 
c: speed of light (3.00  108 m/s)
: wavelength (m, nm, etc.)
: frequency (Hz)
 EX: Find
the frequency of a photon with a
wavelength of 434 nm.
GIVEN:
WORK:
=c
=?

 = 434 nm
= 4.34  10-7 m  = 3.00  108 m/s
-7 m
8
4.34

10
c = 3.00  10 m/s
 = 6.91  1014 Hz
 Planck
(1900)
• Observed - emission of light from hot
objects
• Concluded - energy is
emitted (absorbed or
released) in small, specific
amounts (quanta)
• Quantum - smallest energy packet that can
be emitted or absorbed as EM radiation by
an atom.
Planck proposed that the energy, E, of a single
quantum energy packet equals a constant (h) times its
frequency
The energy of a photon is proportional to its frequency.
E = h
E: energy (J, joules)
h: Planck’s constant (6.6262  10-34
J·s)
: frequency (Hz)
 EX: Find
the energy of a red photon with
a frequency of 4.57  1014 Hz.
GIVEN:
WORK:
E=?
E = h
 = 4.57  1014 Hz
E = (6.6262  10-34 J·s)
h = 6.6262  10-34 J·s
(4.57  1014 Hz)
E = 3.03  10-19 J
 Planck
(1900)
• Observed - emission of light from hot
objects
• Concluded - energy is
emitted (absorbed or
released) in small, specific
amounts (quanta)
• Quantum - smallest energy packet that can
be emitted or absorbed as EM radiation by
an atom.
 Planck
(1900)
vs.
Classical Theory
Quantum Theory


Energy is always emitted or absorbed in whole
number multiples of hv, such as hv, 2 hv, 3 hv, 4hv, ….
The allowed energies are quantized, that is their
values are restricted to certain quantities.
The notion of quantized rather than continuous
energies is strange. Consider a ramp and a staircase,
on a ramp you can vary the length your steps and
energy used on the walk up. When walking up steps
you must exert exactly the specific amount of energy
needed to reach the next step. Your steps on steps are
quantized, you cannot step between them.
 Einstein
(1905)
• Observed - photoelectric effect

Einstein (1905)
 Concluded - light has properties of both
waves and particles (photons)
“wave-particle duality”
 Photon - particle of light that carries a
quantum of energy
 Used planck’s quantum theory to deduced
that: Ephoton = hv
 DEFINITION: A set of frequencies of EM waves emitted by atoms of the element.
EX. neon light absorb electrical energy, e- get excited, become somewhat unstable and
release energy in the form of light, a prism separates the light into an atomic emission
spectrum.
excited state
ENERGY IN
PHOTON OUT
ground state
 e-
exist only in orbits with specific
amounts of energy called energy levels
 Therefore…
• e- can only gain or lose certain amounts of
energy
• only certain photons are produced
65
4
Energy
3
2
1
of photon
depends on the
difference in
energy levels
Bohr’s calculated
energies matched
the IR, visible, and
UV lines for the H
atom
 Each
element has a unique bright-line
emission spectrum.
• “Atomic Fingerprint”
Helium
Bohr’s calculations only worked for
hydrogen! 
 Examples:
• Iron
 Now, we
can calculate for all elements
and their electrons – next section 
 Louis
de Broglie (1924)
• Applied wave-particle theory to e-
• e- exhibit wave properties
EVIDENCE: DIFFRACTION PATTERNS
VISIBLE LIGHT
ELECTRONS
 Heisenberg
Uncertainty Principle
• Impossible to know both the velocity and
position of an electron at the same time
 Schrödinger
Wave Equation (1926)
• finite # of solutions  quantized energy levels
• defines probability of finding an e-
h
 
mv
Wavelength = plank’s constant/mass X velocity
 Orbital
(“electron cloud”)
• Region in space where there is 90% probability
of finding an e-
Orbital
Radial Distribution Curve
 Four
Quantum Numbers:
• Specify the “address” of each electron in an
atom
UPPER LEVEL
1. Principal Quantum Number ( n )
• Main energy level occupied the e• Size of the orbital
• n2 = # of orbitals in
the energy level
2. Angular Momentum Quantum # ( l )
• Energy sublevel
• Shape of the orbital
s
p
d
f
n = # of sublevels per level
n2 = # of orbitals per level
Sublevel sets: 1 s, 3 p, 5 d, 7 f
3. Magnetic Quantum Number ( ml )
• Orientation of orbital around the nucleus
• Specifies the exact orbital
within each sublevel
px
py
pz
 Orbitals
combine to form a spherical
shape.
2px
2py
2s
2pz
4. Spin Quantum Number ( ms )
• Electron spin  +½ or -½
• An orbital can hold 2 electrons that spin in
opposite directions.
 Pauli
Exclusion Principle
• No two electrons in an atom can have the same 4
quantum numbers.
• Each e- has a unique “address”:
1. Principal #
2. Ang. Mom. #
3. Magnetic #
4. Spin #




energy level
sublevel (s,p,d,f)
orientation
electron
Read
Section 4-2!