Ch. 13 Quantum Mechanical
Model
Electron Configuration
Catalyst: Check your homework
answers
1. 28.097 amu
2. 24.2846 amu
3. 35.4529 amu
4. 19.97% Boron-10 80.03% Boron-11
5. 70.88 amu
6. 6.9409 amu
7. 108.9049 amu
8. 19.97% Boron-10 80.03% Boron-11
9. 92.5% Lithium-7 and 7.5% Lithium-6
!2
• After the nucleus was discovered and
studied, scientists continued to develop
the model of the atom and look further
into the electron cloud and the
organization of the electrons
• Neils Bohr first claimed that each electron
had its own fixed energy level and energy
had to be released or gained in order for
electrons to change energy levels
!3
• This led to the quantum mechanical
model of the atom which consists of 4
parts
!4
The Principle Energy Levels (n)
n=1, n=2, n=3, n=4, n=5, n=6, n=7
•
Electrons are in designated energy levels.
Each energy level is subdivided
into
• Sub levels: a collection of orbitals within
an energy level having similar
characteristics (type of room)
• S
• P
• D
• F
• Orbitals:region of space where there is a
high probability of finding an electron
(beds in the room, 1 orbital = 2 electrons)
!6
Each sublevel has orbitals of different shapes
“s” orbital
spherical shaped,
and holds up to
2e-
“p” orbital
Dumbbell shaped
Arranged x, y, z axes, and can
hold up to 6e-
“d” orbital
clover shaped, and can hold up
to 10e-
“f” orbital
• Orbitals combine to
form a spherical shape.
• This orbital can hold up
to
14e-
f
2px
2py
2s
2pz
Rules for e- configurations
1. Aufbau principle: e- enter
orbitals of lowest energy
level
2. Pauli exclusion principle:
an atomic orbital may have
at most 2 e-, e- in the same
orbital will spin in opposite
directions
3. Hund’s rule: when eoccupy orbitals of = energy,
1 enters each orbital until
all the orbitals contain 1 ew/parallel spins.)
N=3
3d ___ ___ ___ ___ ___
(4s ____)
3p ___ ___ ___
3s ___
N=2
2p ___ ___ ___
2s ___
N=1
1s ___
Principle energy level (n)
(Hotel Floors)
Number of different
sublevels on the
floor
Sublevel and Number of Orbitals in each
Sublevel
(Type of room and number of beds in each
room)
1st energy level
1 sublevel
“s” (1 orbital)
2nd
2 sublevels
“s” (1) & “p” (3 orbitals)
3rd
3 sublevels
“s”(1) , “p” (3) & “d” (5 orbitals)
4th
4 sublevels
“s”(1), “p”(3) , “d”(5), and “f” (7)
7s
6s
5s
6p
4f
5p
4d
4p
4s
3s
2s
1s
5d
3p
2p
3d
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
3p
2p
6d 6f 6g
5d 5f 5g
4d 4f
3d
Filling in orbitals then writing the electron
configuration
4p _ ↑↓ _
3d _ ↑↓ _
4s _ ↑↓ _
3p _ ↑↓ _
3s _ ↑↓ _
2p _ ↑↓ _
2s _ ↑↓ _
1s _↑↓_
_ ↑↓ _ _ ↑↓ _
_ ↑↓ _ _ ↑↓ _ _ ↑↓ _ _ ↑↓ _
_ ↑↓ _ _ ↑↓ _
_ ↑↓ _ _ ↑↓ _
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
D. According to their e- configs, elements can be
classified into 4 main groups
1. Noble Gases – outermost s & p sublevels filled
Because they have their s2 & p6
orbitals filled they follow the:
2 + 6 = OCTET RULE
2. Representative Elements – outermost s or p
sublevel is only partially filled, energy level
same as period #
s1
The pink elements excluding the Noble Gases.
s2
p1 p2
p3 p4 p5
3. Transition metals – outermost s sublevel &
nearby d sublevel contain e- , energy level is
the same as the period # minus 1
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10
4. Inner Transition metals - outermost s &
nearby f generally contain e-
f1
f14
Your Periodic Table should look like this.
s1
s2
p1 p2
p3 p4 p5
s2
p6
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10
f1
f14
Learning Check
How many electrons are present in the d sublevel of a
neutral atom of Manganese?
1 2 3 4 5
5 electrons
What element has the electron configuration
1s22s22p63s23p4?
Add together all the exponents, then find that
atomic number. = Sulfur 16
E. Using the Noble Gases to write Shorthand
• Write the noble gas that is in the previous row.
• Use the symbol of the noble gas, put it in
brackets, then write the rest of the
configuration.
Write the e- config using Noble Gas notation for Cobalt.
It would be written [Ar] 4s2 3d7
• Write the e- config for Tin (Sn).
• [Kr] 5s2 4d10 5p2
Learning Check
Using the Noble Gas Shorthand write the e- configuration
1. Cr
[Ar] 4s2 3d4
2. Br
3. Te
[Ar] 4s2 3d10 4p5
4. Ba
[Kr] 5s2 4d10 5p4
[Xe] 6s2
Electromagnetic Spectrum
• The electromagnetic spectrum (see p. 373) includes
radio waves, microwaves, infrared waves, visible light,
ultraviolet waves, x-rays, and gamma rays.
• Visible light is in the middle of the spectrum.
• The speed of light is 3.0 X 108 m/s.
• The formula for light is
c =ƛ
• C = speed of light, ƛ = wavelength,
= frequency
– Visible light has many wavelengths of light that can be
separated into red, orange, yellow, green, blue,
indigo, and violet (ROY G BIV)
Atomic Emission Spectrum
• Every element gives off light when it is
excited by the passage of an electric current
through its gas or vapor.
• The atomic emission spectrum occurs when
the light that is given off by an element in its
excited state is passed through a prism. It
consists of a few lines called a line spectra or
discontinuous spectra. Each line on the
spectra corresponds with a frequency.
• See page 374.
• Work problems # 11 and 12 on page 375.
Planck’s Constant
• In 1900, German Physicist Max Planck used
math to explain why objects, such as iron,
that are heated change color.
• He said energy can be quantized. The size of
an emitted or absorbed quantum depends on
the size of the energy change. A small energy
change involves the emission or absorption of
low frequency radiation. A large energy
change involves the emission or absorption of
high frequency radiation.
Planck’s constant cont.
• In 1905, Albert Einstein used Planck’s
work to call quanta of light photons. He
then used this information to explain the
photoelectric effect (metals eject/emit
electrons called photoelectrons when
light shines on them).
• Work problems 13 and 14 on p. 379.
Planck’s constant cont.
• The math formula used is:
E=hxv
E = radiant energy of a unit (quantum)
h = Planck’s constant = 6.6262 x 10 -34
v = frequency of radiation