Transcript Electron Notes

Ch. 5 - Electrons in Atoms

Wavelength () - length of one complete wave
measured in m, cm, or nm


Frequency () - # of waves that pass a point during
a certain time period,


In light it tells us which color it is
hertz (Hz) = 1/s
Amplitude (A) - distance from the origin to the
trough or crest

how much energy the wave is carrying. It is the height of
the wave. It is measured in meters. In SOUND it tells us
how LOUD it is. In LIGHT it tells how BRIGHT it is.

crest
A
greater
amplitude
origin
A

trough
greater
frequency
• 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
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)
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
 Dispersed light falls on metal samples, the
different frequencies produce different
energetic photoelectrons

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
Ch. 5 - Electrons in Atoms
Set of frequencies of EM waves emitted by atoms an element
when they absorb electrical energy, eˉ get excited, become
somewhat unstable and release energy in the form of light
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

Ground state: lowest allowable atomic electron
energy state

Excited state: any higher energy state
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.


Helium
Examples:


“Atomic Fingerprint”
Iron
Now, we can calculate for all elements and their electrons
Ch. 5 - Electrons in Atoms

Louis de Broglie (1924)

Proposed eˉ in their orbits behave like a wave

Wavelength of an eˉ depends on its mass(m)
and its velocity (v):
λ = _h _
mv
EVIDENCE: DIFFRACTION PATTERNS
VISIBLE LIGHT
ELECTRONS

Heisenberg Uncertainty Principle

Impossible to know both the velocity and position of an
electron at the same time

Attempting to observe an electron’s position changes its
momentum and attempting to observe an electron’s
momentum changes its position. Therefore electrons cannot
be locked into well-defined circular orbits around the
nucleus.



Schrödinger Wave Equation (1926)
proposed a wave equation incorporating both the
wave and particle nature of the electron.
The result of the equation, wave functions, shows the
probability that an electron will be in a certain region
of space at a given instant. This electron density is
represented by a distribution of dots which
represents where electrons are located about 90% of
the time

finite # of solutions  quantized energy levels

defines probability of finding an e-



Orbital (“electron cloud”)
a specific distribution of electron density in space.
Each orbital has a characteristic energy and shape.
Orbital
Specify the “address” of each
electron in an atom
UPPER LEVEL
1. Principal Quantum Number (n = 1, 2, 3, …)
(see periodic table left column)

Indicates the relative size and energy of
atomic orbitals

As (n) increases, the orbital becomes larger, the
electron spends more time farther from the
nucleus

Each major energy level is called a principle
energy level
Ex: lowest level = 1 ground state,
highest level = 7 excited state
2. Energy Sublevel

Defines the shape of the orbital (s, p, d, f)

# of orbital related to each sublevel is always an odd #
s = 1, p = 3, d = 5, f = 7

Each orbital can contain at most 2 electrons
s
p
d
f
Subscripts x, y, z designates orientation

Specifies the exact orbital within each sublevel
px
py
pz
4. Spin Quantum Number ( ms )

Electron spin  +½ or -½

An orbital can hold 2 electrons that spin in opposite
directions.

Pauli Exclusion Principle

A maximum of 2 electrons can occupy a single
atomic orbital

Only if they have opposite spins
1. Principal #
2. Energy sublevel
3. Orientation
4. Spin #




energy level
(s,p,d,f)
x, y, z
exact electron
Ch. 5 - Electrons in Atoms
IV. Electron
Configuration
A. General Rules
Aufbau Principle
Electrons fill the
lowest energy
orbitals first.
“Lazy Tenant
Rule”
A. General Rules
Hund’s Rule
Within a sublevel, place one e- per
orbital before pairing them.
“Empty Bus Seat Rule”
WRONG
RIGHT
Notation
s
p
1
2
3
4
5
6
7
f(n-2)
d (n-1)
6
7
© 1998 by Harcourt Brace & Company
B. Notation
Orbital Diagram
O
8e-
1s
2s
Electron Configuration
2
2
4
1s 2s 2p
2p
B. Notation
Longhand Configuration
S 16e- 1s2 2s2 2p6 3s2 3p4
Core Electrons
Valence Electrons
Valence electrons: determine chemical properties of that
element & are the electrons in the atoms outermost orbital
Shorthand Configuration
S 16e- [Ne] 3s2 3p4
Shorthand Notation
Shorthand Configuration
Core e-: Go up one row and over to the
Noble Gas.
Valence e-: On the next row, fill in the
# of e- in each sublevel.
1
2
3
4
5
6
7
C. Periodic Patterns
Example - Germanium
1
2
3
4
5
6
7
[Ar]
2
4s
10
3d
2
4p
D. Stability
Full energy level
Full sublevel (s, p, d, f)
Half-full sublevel
1
2
3
4
5
6
7
D. Stability
Electron Configuration Exceptions
Copper
EXPECT:
[Ar] 4s2 3d9
ACTUALLY:
[Ar] 4s1 3d10
Copper gains stability with a full
d-sublevel.
D. Stability
Electron Configuration Exceptions
Chromium
EXPECT:
[Ar] 4s2 3d4
ACTUALLY:
[Ar] 4s1 3d5
Chromium gains stability with a half-full
d-sublevel.
D. Stability
Ion Formation
Atoms gain or lose electrons to become
more stable.
Isoelectronic with the Noble Gases.
1
2
3
4
5
6
7
D. Stability
Ion Electron Configuration
Write the e- config for the closest Noble
Gas
EX: Oxygen ion  O2-  Ne
2O
10e
[He]
2
2s
6
2p
Read
Section 5-3!