Transcript Sections 1 – 3 (Read pages 126-153)

Chapter 5: Electrons in Atoms
S E CTI ONS 1 – 3 (R E A D P A GE S 1 2 6 - 1 5 3)
Chapter 5 Vocabulary
• Amplitude
• Heisenburg uncertainty principle
• Atomic emission spectrum
• Hertz
• Atomic orbital
• Hund’s Rule
• Aufbau principle
• Pauli Exclusion Principle
• Electromagnetic radiation
• Photons
• Electron configurations
• Quantum
• Energy levels
• Quantum mechanical model
• Frequency
• Spectrum
• Ground state
• Wavelength
Section 5.1: Models of the Atom
R E A D P A GE S 1 2 7 - 1 3 2
Rutherford Planetary Model (1911)
• Based on the idea that protons and neutrons
existed in the nucleus and electrons moved
around the nucleus, much like planets around
the sun
• Explained simple properties of the atom
• Could not explain the chemical properties of
elements
Bohr Model (1913)
• Neils Bohr
• Danish physicist
• Considered Hydrogen because it is
the simplest atom with only 1 electron
• Proposed that an electron is found
only in specific circular paths, or
orbits, around the nucleus of the atom
• Each orbit has a fixed energy value –
and is referred to as an energy level
Think of the energy levels
like a ladder
• Each rung represents an energy level
• The lowest energy level is the lowest rung
• Just like a person can climb a ladder from rung
to rung, electrons can jump from energy level
to energy level
• A person cannot stand between rungs, and
electrons cannot exist between energy levels
• A person must move a specific distance from
rung to rung, electrons must gain or lose a
certain amount of energy to move from energy
level to energy level
Energy
• Quantum of energy
• The amount of energy required to move an
electron from one energy level to another
energy level
• The amount of energy an electron gains or
loses is not always the same
The problem with the Bohr Model
• While it explained the hydrogen atom, it
failed in many ways to explain energy
changes in atoms with more than one
electron
Quantum Mechanical Model (1926)
• Determines the allowed energies that an
electron can have (similar to Bohr Model)
• Determines how likely it is to find the electron
in various locations around the nucleus (not
an exact path)
• Referred to as electron cloud
Atomic Orbitals and Sublevels
• A region of space with a high probability of finding an electron
• Energy level
• Defined by principal quantum numbers (n)
• n = 1 is the first principal energy level
• n = 2 is the second principal energy level
• Up through n = 7
• For each principal energy level, there may be several orbitals with
different shapes and different energy values
• Called sublevels
Sublevels
• Four Types
1.
2.
3.
4.
S
P
D
F
S Sublevel
• Shaped like a sphere
• Contains 1 orbital
• Can hold 2 electrons
• Exists at all principal energy level values
(n = 1 through n = 7)
P Sublevel
• Each orbital shaped like a dumbbell
• Contains 3 orbitals
• Can hold 6 electrons
• Exists at principal energy levels n = 2
through n = 7
D Sublevel
• Contains 5 orbitals
• Can hold 10 electrons
• Exists at principal energy levels n = 3
through n = 6
F Sublevel
• Contains 7 orbitals
• Can hold 14 electrons
• Exists at principal energy levels n = 4 and n = 5
Sublevel Summary Chart
Sublevel
Shape
S
P
D
-
F
-
Number of Number of
Orbitals
Electrons
Principal
Energy
Levels
Sublevel Summary Chart
Sublevel
Shape
Number of Number of
Orbitals
Electrons
Principal
Energy
Levels
S
Sphere
1
2
1–7
P
Dumbbell
3
6
2–7
D
-
5
10
3–6
F
-
7
14
4 and 5
Quick Quiz
S E CTI ON 5 . 1
1. How many possible sublevels are
there in the 2nd principal energy level?
1. One
2. Two
3. Three
4. Four
2. How many electrons can a 4f sublevel
hold?
1. Two
2. Six
3. Ten
4. Fourteen
3. What possible sublevels exist for n = 3?
1. S
2. S and P
3. S, P, and D
4. S, P, D, and F
What’s Next?
• Book Work
• Page 132 #1-7
Page 132 #1-7
1. Why did Rutherford’s atomic model need to be replaced?
2. What was the basic new proposal in the Bohr model of the atom?
3. What does the quantum mechanical model determine about electrons in
atoms?
4. How do two sublevels of the same principal energy level differ from each other?
5. How can electrons in an atom move from one energy level to another?
6. The energies of electrons are said to be quantized. Explain what this means.
7. How many orbitals are in the following sublevels?
a. 3p sublevel
b. 2s sublevel
c. 4p sublevel
d. 3d sublevel
e. 4f sublevel
#1 – Why did Rutherford’s atomic model
need to be replaced?
•It could not explain the
chemical properties of the
elements.
#2 – What was the basic new proposal in
the Bohr model of the atom?
•An electron is found in
specific paths, called orbits,
around the nucleus
#3 – What does the quantum mechanical
model determine about electrons in atoms?
• The allowed energies an
electron can have
• How likely it is to find an
electron in various locations
#4 – How do two sublevels of the same
principal energy level differ from each other?
• Sublevels have different
shapes
#5 – How can electrons in an atom
move from one energy level to another?
• Electrons can jump energy
levels by gaining or losing
specific amounts of energy
#6 – The energies of electrons are said to be
quantized. Explain what this means.
• Fixed amounts of energy to
move between energy levels
• Not always the same
amount
#7 – How many orbitals are in the
following sublevels?
a.3p sublevel
b.2s sublevel
c.4p sublevel
d.3d sublevel
e.4f sublevel
Section 5.2 – Electron Arrangement in Atoms
R E A D P A GE S 1 3 3 - 1 3 6
Electron Configurations
• Electrons interact with the nucleus in such a way that they
make the most stable arrangement.
• Called the electron configuration
• Every element on the periodic table has a unique electron
configuration
Aufbau Principle
• Electrons occupy the orbitals of lowest energy first
Pauli Exclusion Principle
• Atomic orbitals can
hold up to 2 electrons,
which have opposite
spin
Hund’s Rule
If you were told to draw a 3p orbital
with 3 electrons, you could draw it
multiple ways
1
2
3
• Electrons occupy
orbitals in the same
energy level in a way
that maximizes the
number of electrons
with the same spin
• Rule: Add electrons
individually, then pair up
More with Electron Configurations
2
1s
Principal Energy
Level (n)
Sublevel
Number of
electrons
Let’s Try It!
Write the electron
configuration for oxygen.
Practice Problems
• Write the electron configurations for the
following elements:
A. Sodium
B. Chlorine
C. Krypton
D. Aluminum
Electron Configurations Continued
Write the electron configuration for Barium:
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2
Wouldn’t it be nice if we could make that shorter?
More electron configuration practice
• Write the electron configurations for the
following elements:
A. Potassium
B. Carbon
C. Beryllium
D. Selenium
Abbreviated Electron Configurations
• Using the noble gases, we can abbreviate electron
configurations
• Here’s how it works:
• Find the element you need on the periodic table.
• Move up 1 row and to the right until you reach the noble gas
column.
• Write the chemical symbol of the noble gas in brackets [ ].
• Continue the electron configuration from there.
Abbreviated Example
Follow the
steps:
• Find the element you
need on the periodic
table.
• Move up 1 row and to the
right until you reach the
noble gas column.
• Write the chemical
symbol of the noble gas
in brackets [ ].
• Continue the electron
configuration from there.
[Xe]
2
6s
Another Abbreviated Example - Bromine
2
5
10
[Ar] 4s 3d 4p
Stability in Electron Configurations
• Atoms want to have full s and p sublevels to be stable
• Stability can also be achieved through a half filled sublevel
• S sublevels hold 2 electrons maximum, and are stable with 1
or 2 electrons
• P sublevels hold 6 electrons maximum, and are stable with 6
or 3 electrons
Electron Configuration Exceptions
• Copper and Chromium are two exceptions to the
Aufbau principle
• Look at the 4s and 3d sublevels
• Predicted configurations
• Cu – 1s2 2s2 2p6 3s2 3p6 4s2 3d9
• Cr – 1s2 2s2 2p6 3s2 3p6 4s2 3d4
• Actual configurations
• Cu – 1s2 2s2 2p6 3s2 3p6 4s1 3d10
• Cr – 1s2 2s2 2p6 3s2 3p6 4s1 3d5
Quick Quiz
S E CTI ON 5 . 2
1. Which statement correctly defines the
Aufbau principle?
1. Electrons occupy orbitals of the same energy in a
way that maximizes the number of same spin
directions
2. Orbitals can hold 2 electrons maximum
3. Electrons occupy the orbitals of lowest energy first
4. Orbitals can hold up to 8 electrons
2. What is the correct electron
configuration for potassium?
2
2
6
1.1s 2s 2p
2.1s2 2s2 2p6 3s1
2
2
6
2
3.1s 2s 2p 3s
4.1s2 2s2 2p6 3s2 3p1
3. Identify the element that has the following
electron configuration:
1s2 2s2 2p6 3s2 3p3
1. Al
2. Si
3. P
4. S
What’s Next?
• Book Work
• Page 136 #10 and 11
Page 136 #10-11
10.What are the three rules for writing
electron configurations of elements?
11.Explain why the actual electron
configurations for some elements differ
from those assigned using the Aufbau
principle.
#10 – What are the three rules for writing
electron configurations of elements?
•Aufbau Principle - Electrons occupy the
orbitals of lowest energy first
•Pauli Exclusion Principle - Atomic orbitals
can hold up to 2 electrons, which have
opposite spin
•Hund’s Rule - Add electrons individually,
then pair up
#11 – Explain why the actual electron
configurations for some elements differ from
those assigned using the Aufbau principle.
• Half filled sublevels are not as stable a
full sublevels, but are more stable than
other configurations
Homework: Electron Configuration Problems
1. Write the full electron configurations for:
1. Magnesium (Mg)
2. Rubidium (Rb)
3. Nitrogen (N)
2. Write the abbreviated configurations for:
1. Sulfur (S)
2. Iodine (I)
3. Zinc (Zn)
Section 5.3 – Physics and the Quantum
Mechanical Model
R E A D P A GE S 1 3 8 - 1 4 5
Light
• Isaac Newton thought light consisted of particles
• By 1900, there was enough experimental evidence
to prove that light actually consists of waves
Parts of a Wave
• Crest
• Highest point of a
wave
• Trough
• Lowest point of a
wave
• Amplitude
• Wave’s height from
zero to the crest or
trough
• Wavelength ()
• The distance between two crests or
troughs
• Symbol is Greek letter Lambda
• Frequency ()
• The number of wave cycles that
passes a given point in a certain
amount of time
• Symbol is Greek letter Nu
• Unit = Cycles/Second = Hertz (Hz)
Labeling the Parts of a Wave
Crest
Wavelength
Amplitude
Trough
64
Wavelength and frequency are inversely
proportional to each other, as the
wavelength of light increases, the
frequency will decrease.
65
Speed of Light
• Formula
• c = 
• c = constant, speed of light (2.998 x 108 m/s)
•  = wavelength
•  = frequency
66
1. What is the wavelength of radiation with a frequency of 1.50 x 1013 Hertz?
2. What frequency is radiation with a wavelength of 5.00 x 10-8 m?
1.  = c = 2.998 x 108 m/s

1.5 x 1013 Hz
= 1.998 x 10-5 m
2.  = c = 2.998 x 108 m/s

5.00 x 10-8 m
= 5.996 x 1015 s-1
Electromagnetic Spectrum
• Displays a range of wavelengths
• Electromagnetic radiation includes radio
waves, microwaves, infrared waves, visible
light (the rainbow), ultraviolet waves, X-rays,
and gamma rays.
68
69
Atomic Spectra
• Ground state
• Electrons in their lowest energy level
• Excited state
• When atoms absorb energy, electrons move into a
higher energy level
• When electrons return to their lower energy
level, they emit light, each having their own
color depending on the amount of energy
released
70
Energy of an Electron
• The light emitted by an electron
moving from a higher to lower energy
level has a certain amount of energy
•E = h x v
• E = Energy in J
• h = Planck’s constant (6.626 x 10-34 J-s)
• v = frequency in Hertz
71
Series of Light
(Specific Electron Transitions in Hydrogen Atom)
• Lyman Series
• Transitions that end on n = 1
• Ultraviolet light
• Balmer Series
• Transitions that end on n = 2
• Visible light
• Paschen Series
• Transitions that end on n = 3
• Infrared light
72
Back to the idea of light as particles…
• Einstein explained that light could be
described as quanta of energy called
photons
• Quanta behave as particles
73
Quantum Mechanics
• Mechanics describes motion
• Quantum mechanics is the motion of
subatomic particles and atoms as waves
• Heisenberg Uncertainty Principle
• It is impossible to know exactly the velocity (speed)
and position of a particle at the same time
• Only applies to subatomic particles
74
Quick Quiz
S E CTI ON 5 . 3
75
1. Calculate the frequency of a radar wave with a
wavelength of 125 m.
a. 2.40 x 109 Hz
b. 2.40 x 1024 Hz
c. 2.40 x 106 Hz
d. 2.40 x 102 Hz
76
2. The lines in the emission spectrum for an element are
caused by
a. the movement of electrons from lower to higher
energy levels.
b. the movement of electrons from higher to lower
energy levels.
c. the electron configuration in the ground state.
d. the electron configuration of an atom.
77
What’s Next?
• Book Work:
• Page 146 #16-21
78
Page 146 #16-21
16.How are wavelength and frequency related?
17.Describe the cause of atomic emission spectrum of an element.
18.How is the change in electron energy related to the frequency of light
emitted in atomic transitions?
19.How does quantum mechanics differ from classical mechanics?
20.The lines at the ultraviolet end of the hydrogen spectrum are known as
the Lyman series. Which electron transitions within an atom are
responsible for these lines?
21.Arrange the following in order of decreasing wavelength:
a. infrared radiation from a heat lamp
b. dental x-rays
c. signal from a shortwave radio station
#16 – How are wavelength and
frequency related?
•Inversely proportional – as one
increases, the other decreases
#17 – Describe the cause of atomic
emission spectrum of an element.
•Atoms absorb energy and
electrons move to higher energy
levels. When the electrons fall
back to their ground states, they
release energy in the form of light
#18 – How is the change in electron energy
related to the frequency of light emitted in
atomic transitions?
• The light emitted by an electron
moving to a lower energy level has
a frequency directly proportional
to the energy change of the
electron.
#19 – How does quantum mechanics
differ from classical mechanics?
• Classical mechanics – describes
the motion of large bodies
• Quantum mechanics – describes
the motion of subatomic particles
and atoms as waves
#20 – The lines at the ultraviolet end of the hydrogen
spectrum are known as the Lyman series. Which electron
transitions within an atom are responsible for these lines?
• All transitions ending on n = 1
#21 – Arrange the following in order of decreasing wavelength:
a. infrared radiation from a heat lamp
b. dental x-rays
c. signal from a shortwave radio station
• C>A>B