Download 1,0-,1,2 + ½

Chapter 5 : Electrons in Atoms
Problems with Rutherford’s
Model of the Atom
Chlorine # 17
Argon # 18
Potassium # 19
Reactive
Not reactive
Very reactive
Does not explain why elements react the way they do!
• It’s because they have
different arrangements of
electrons!
Electrons and Light
• An element’s chemical behavior is related to
its arrangement of electrons
• Elements give off different light when burned
• We can analyze light to learn about atoms!
Atomic Emission Spectra
The frequencies of the electromagnetic radiation (EMR) emitted
by atoms of an element are unique to each element, like
fingerprints on people!
The Quest for a Better Model
• Electromagnetic radiation (EMR) behaves
like a wave.
Characteristics of a Wave
Wavelength = λ
Frequency = f
(number of waves that pass a point per second)
1 Hertz (Hz) = 1 wave per second (SI Unit for frequency)
Types of Waves
Compressional or Longitudinal Wave (sound or slinky)
Transverse Wave (light, rope)
://www.youtube.com/watch?v=Rbuhdo0A
ZDU
Light as a Wave
• We can relate the speed, frequency, and
wavelength of EMR with the equation:
v = fl
v is velocity (m/s)
f is frequency (Hertz – Hz, 1 Hz = 1 wave/sec)
l is wavelength (m)
The speed of light (c) is a constant!
c = 3.0 x 108 m/s
Velocity and Frequency of Light
v = λf
v (or c for light) = speed of light (3.0 x 108 m/s)
↑ wavelength ↓ frequency
↓ wave length ↑frequency
Problems
A helium-neon laser emits a light with a wavelength of
633 nm. What is the frequency of this light?
Use the equation v = λf
First convert nm to m.
633 nm X
1 m
1 x 109
nm
= 6.33 x 10-7 m
Then do the math.
3.0 x 108 m/s = (6.33 x 10-7 m) (f)
(6.33 x 10-7 m)
(6.33 x 10-7 m)
4.7 x 1015 1/s or Hz = f
An FM radio station broadcasts at a frequency of 98.5
MHz. What is the wavelength of the station’s
broadcast signal?
Frequency = 98.5 MHz
Velocity = 3 x 108 m/s
Wavelength = X
Use the formula,
Rearrange the formula,
Convert 98.5 MHz to Hz
98.5 MHz 1 x 106 Hz = 9.85 x 107 Hz
1MHz
λ = vf so plugging in our values:
λ = 3 x 108 m/s / 9.85 x 107 Hz = 3.05 m
v = λf
λ = v/f
Stop.
Light: Particle or Wave?
Wave model doesn’t
address:
Why heated objects emit
only certain frequencies of
light at a given temperature?
Why some metals emit
electrons when a colored
light of a specific frequency
shines on them?
(photoelectric effect)
http://www.youtube.com/watch?v=WO38qVDGgqw&feature=related
Iron
Dark gray = room temp
Red = hot temp
Blue = extremely hot temp
Photoelectric Effect –
The problem with wave theory
Only certain frequencies of light could emit an electron
from a plate of Ag.
Accumulation of low frequencies couldn’t
Einstein and the Dual
Nature of EMR (1900)
• Electromagnetic Radiation
(EMR) acts as a wave of
individual particles (photons)
Wave vs. Particle
Wave Characteristics
Particle
Characteristics
Light can be reflected.
Photons have the ability
to knock individual
electrons off of a
conductor-Photoelectric
Effect!
Light shows interference Metals glow only specific
patterns.
colors when heated.
Wave
+
Wave
RESULT
What is the relationship between
energy and frequency?
Max Planck - 1900
Matter gains or loses energy only in small, specific
amounts called quanta. This is why we see specific
color lines in emission spectra.
A quantum is the minimum amount of energy that
can be gained or lost by an atom
Equantum = hf
E is energy
h is Planck’s constant = 6.626 x 10-34 J·s
f is frequency
J is joule, SI Unit for energy
Calculating the energy in a Photon
Ephoton = hf
E = (6.626 x 10-34 J·s) x (7.23 x 1014 s-1)
E = 4.79 x 10-19 J
Niels Bohr - 1913
• Worked in Rutherford’s lab
• Proposed a quantum model of the atom
• Explains why emission spectra were
discontinuous
• Predicted frequencies of light in Hydrogen’s
atomic emission spectra
Bohr’s Explanation
• Ground state – lowest energy state of an atom’s electrons
• Excited state – when an atom’s electrons gain energy
• Electrons move in circular orbits
– Smaller orbit – lower energy state, “energy level”
– Larger orbit – higher energy state, “energy level”
An explanation for Emission Spectra
http://www.mhhe.com/phys
sci/astronomy/applets/Bohr/
applet_files/Bohr.html
Excited States
Ground state
Atoms absorb energy
and are excited. As the
electron returns to the
ground state they give off
energy “photon” equal to
the difference in energy
levels.
Bohr’s Model
Electrons move around the nucleus in set
orbits in specific energy levels. When
excited, electrons give off a discrete amount
of energy as an emission spectrum.
This discrete amount of energy is a quantum.
Click the link to make an electron jump!
http://www.mhhe.com/physsci/astronomy/applets/Bohr/applet_files/Bohr.html
Bohr used work of others…
• Balmer—made an equation (math) to
connect the lines of the hydrogen
spectrum to each other.
• Planck—Energy is directly proportional to
the frequency of light.
Problem: Bohr’s Model Only
explains Hydrogen
• Louis de Broglie (1924) – proposed that
the energy levels are based on the wave
like nature of electrons
Heisenberg Uncertainty Principle
• It is impossible to know the velocity and position
of an electron at any given time
• Bohr gives a specific place for electrons, while
Heisenberg says you can’t know where the
electrons are at a particular time.
Photon and
electron are
about the
same mass.
Erwin Schrodinger - 1926
• Developed the quantum mechanical model
of the atom
– Assigns electrons to energy levels like Bohr
– Does not predict the path of the electron
– It predicts the probability of finding an electron
• An electron’s “atomic orbital” is the most probable
location of the electron at any point in time.
Each dot is a picture
of an electron during
a given amount of
time.
Where does the
electron spend most
of the time?
Boundary
represents the
location of an
electron 90% of the
time.
Stop.
Each dot is a picture
of an electron during
a given amount of
time.
Where does the
electron spend most
of the time?
Boundary
represents the
location of an
electron 90% of the
time.
Principle Energy Levels
• 7 energy levels
• Lowest energy is 1 – greatest energy 7
• Each level consists of sublevels
The second energy level is
larger and the electrons are
farther from the nucleus.
s Orbitals
p and d Degenerate Orbitals
Degenerate orbitals have exactly the same amount of energy.
f Degenerate Orbitals
Quantum
Number
Symbol
Possible
Values
Definition
Principle
n
1,2,3…
Major energy
Level or shell
Angular
Momentum
l
0 to n-1
Sublevel or
subshell
Magnetic
ml
-l to +l
Orbital
orientation
Spin
ms
+½ to -½
Spin direction
If Principle
(n) =
Then Angular
(l) =
And Magnetic
(ml) =
1
0 (s)
0
2
0 (s)
1 (p)
0
-1, 0, +1
3
0 (s)
1 (p)
2 (d)
0 (s)
1 (p)
2 (d)
3 (f)
0
-1, 0, +1
-2,-1,0,+1,+2
0
-1, 0, +1
-2,-1,0,+1,+2
-3,-2,-1,0,1,2,3
4
0 = s, 1 = p, 2 = d, 3 = f
To Write Ground State Electron
Configurations…
1. Lowest energy is the most stable
2. 3 principals or rules to follow—Pauli
Exclusion Principle, Aufbau Principle,
and Hund’s Rule
Pauli Exclusion Principle
For electrons to occupy the same orbital,
they must have opposite spin.
• That limits 2 electrons per orbital, written
as up or down:
Aufbau Principle – each electron must
occupy the lowest energy state possible.
What that means is:
1. Within a principle energy level, the energy sublevels
have different energies.
2. The sublevels increase in energy from s,p,d,f
3. Orbitals in a given energy sublevel have equal
energy (they are degenerate).
4. Principal energy levels can overlap.
Hund’s Rule
In a given sublevel, electrons must occupy
each degenerate orbital before additional
electrons can be added.
Pauli Exclusion Principle
Aufbau Diagram
Tells you the order orbitals are filled with
electrons, in order of increasing energy!
4f
3d 4d 5d
2p 3p 4p 5p 6p
1s 2s 3s 4s 5s 6s 7s
5f
6d
7p
8s
14 e- max
10 e- max
6 e- max
2e- max
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, etc.
How can we write this easily?
Only list the total electrons in each orbital, as a superscript
1. H – 1s1
2. He –1s2
3. Li - 1s2, 2s1
4. Be - 1s2, 2s2
5. B - 1s2, 2s2, 2p1
6. C - 1s2, 2s2, 2p2
7. N - 1s2, 2s2, 2p3
8. O - 1s2, 2s2, 2p4
9. F - 1s2, 2s2, 2p5
10.Ne - 1s2, 2s2, 2p6
Orbital Notation
• Very similar to electron configuration, only
this one takes spin of the electron into
account and may include all of the
electrons or just the valence (outer energy
level) electrons
sublevel
Symbol
Be
Energy Level
1S
2S
↑↓
↑↓
+ and - spin
2Px 2Py 2Pz
Representing Electron
Configurations
Electron Configurations
Sub level diagram – indicates the order that
orbitals are filled
What are the orbital diagrams and electron
configuration notation for Al and Cl?
Electron Configuration Shorthand
• Substitute noble gases from preceding
energy levels in the notation
Li – [He] 2s1
C – [He] 2s2 2p2
Stop.
Valence Electrons
• Electrons in the outer most energy levels
S: [Ne] 3s2 3p4
Sulfur has 6 valence electrons (add 2 from s
and 4 from p)
How many valence electrons do Ne, Al, and
Cl have?
Valence Electrons
Ne: 1s2, 2s2, 2p6
8 e- on 2nd energy level
Al: 1s2, 2s2, 2p6, 3s2, 3p1
3 e- on 3rd energy level
Cl: 1s2, 2s2, 2p6, 3s2, 3p5
7 e- on 3rd energy level
Lewis Dot Models
• Only valence electrons – outermost
electrons (highest energy level)
• Octet rule – all atoms want to have 8
electrons (H and He want 2) in outer orbit
1 2
5
8
X
7 4
Symbol
3
6
Writing Electron Dot Structures
• Fill the valence electrons 1 at a time in any
particular order.
*
Ca *
*
C
* *
*
*O**
* *
*
What are the electron dot diagrams for K, Ar and F?
Electron Dot Structures
Valence
electrons are
used in reactions
and are
represented by
an electron dot
structure.
O
H
Ne
N
Periodic Table Blocks
• Where an element is on the periodic table can
give you clues about its electron configuration.
• For the representative elements (A groups), the
A-group number tells you how many valence
electrons!
– Alkali earth metals (2A*) have 2 valence e– Halogens (7A*) have 7 valence e*Note: A different method of naming the groups numbers the
columns 1-13 starting on the left side of the table and includes
the transition metals. In this system group 2A = group 2,
group 3A = group 13
END OF CHEM 1 UNIT 5!!!
(the rest is for Pre-AP)
Modern Atomic Theory
• Any electron in an atom can be described
by 4 quantum numbers
• Principal Quantum Number
• Azimuthal (Angular Momentum) Quantum
Number
• Magnetic Quantum Number
• Spin Quantum Number
Principal Quantum Number (n)
Related to the size and energy of principal
energy level.
The farther away from the nucleus the more
energy the electron has
1 < 2 < 3 < 4 < 5 < 6 etc….
Azimuthal Quantum Number
(Angular Momentum) = l
• Refers to the
subshells in each
principal energy level
(n)
n
1
2
3
•
•
•
•
S=0
P=1
D=2
F=3
4
l
0
0
1
0
1
2
0
1
2
3
Magnetic Quantum Number (ml)
• Specifies
the orbital
within a
energy
level
where an
electron is
likely to be
found
n
l
Orbital
designation
ml
1
0
1s
0
2
0
2s
0
1
2p
-1,0,+1
0
3s
0
1
3p
-1,0,+1
2
3d
-2,-1,0-,1,2
0
4s
0
1
4p
-1,0,+1
2
4d
-2,-1,0-,1,2
3
4f
-3,-2,-1,0,1,2,3
3
4
Spin Quantum Number (ms)
• + ½ or – ½
• Electrons in the same orbitals must have
opposite spins (Pauli Exclusion Principle)
n
l
Orbital
designation
ml
ms
1
0
1s
0
+ ½, - ½
2
0
2s
0
+ ½, - ½
1
2p
-1,0,+1
+ ½, - ½
0
3s
0
+ ½, - ½
1
3p
-1,0,+1
+ ½, - ½
2
3d
-2,-1,0-,1,2
+ ½, - ½
0
4s
0
+ ½, - ½
1
4p
-1,0,+1
+ ½, - ½
2
4d
-2,-1,0-,1,2
+ ½, - ½
3
4f
-3,-2,-1,0,1,2,3
+ ½, - ½
3
4
A
n
l
ml
ms
2
1
-1
+ ½ or
What are the quantum numbers for A? B?
2,1,0,-½
3,0,0,+½
B
So to summarize
There are four quantum numbers that
describe every electron
– Principal Quantum # - Energy Level (1-7)
– Angular Momentum or Azimuthal Quantum # Type of Orbital (s,p,d,f)
– Magnetic Quantum # - Orientation of the
orbital (x, y, z, xy, xz, yz, etc.)
– Spin Quantum # - Spin (+ or -)