Transcript energy levels

Electrons in Atoms
Ch. 13
Models of the Atom
13-1
The evolution of Atomic Models
• Dalton (1766-1844): atom indivisible
• J.J. Thomson (1856-1940):
– “Plum-pudding” model – negative electrons
stuck in positively charged material
• Rutherford (1871-1937):
– Electrons surround dense nucleus, rest of
atom is empty space
• Bohr (1885-1962):
– “Planetary model,” electrons fixed in energy
levels around nucleus
The Quantum Mechanical
Model
• Quantum Mechanical Model:
– Estimates the probability of finding an
electron in a certain place using the
Schrodinger equation.
– “Fuzzy cloud” model; where the
cloud is more dense the probability of
finding the electron is high, where the
cloud is less dense the probability is
low.
Energy Levels
• Electrons move around the
nucleus in energy levels.
• Quantum of energy = amount of E
required to move to a higher level.
• When they move towards the
nucleus (down a level) they
release energy
• When they move away from the
nucleus (up a level) they require
energy.
• The farther from the nucleus the
energy level is, the more energy is
required to move up a level (away
from nucleus).
Energy Levels
• Quantum number (n) refers to an energy level
– n = 1, 2, 3, 4, …7, values increase going away
from nucleus.
– Each energy level fits a certain amount of
electrons:
• Level 1 = 2 electrons
• Level 2 = 8 electrons
• Level 3 = 18 electrons
• Level 4 = 32 electrons
Sublevels + Orbitals
• Within each energy level there are sublevels; the
number of sublevels is equal to the quantum
number.
– Ex: Energy level 4 has 4 sublevels within it.
• A sublevel is made up of atomic orbitals: s, p, d, f
– Orbital s fits 2 electrons total
– Orbital p fits 6 electrons total
– Orbital d fits 10 electrons total
– Orbital f fits 14 electrons total
– s fills up first, then p, then d, then f
Energy
Level (n)
Sublevel/Orbi
tal
Electrons in
each
Sublevel/Orbita
l
Total # of
electrons in
Level
1
1s
2
2
2
2s
2p
3s
3p
3d
2
6
2
6
10
18
4s
4p
4d
4f
2
6
10
14
32
3
4
8
Atomic Orbitals
Orbital
Shape
# of Electrons
s
Spherical
2
p
Dumbbell
6
d
Clover-leaf
10
f
Complex
14
f - orbital
s - orbital
d - orbital
DRAW!
p - orbital
Electron Arrangement in
Atoms
13-2
Label + Color on white, blank, large periodic table!
d-1
5d
1
6d
1
f-2
Electron
Arrangement in Atoms
•
•
•
•
•
•
•
•
Period 1 - 1s2
Period 2 - 2s22p6
Period 3 - 3s23p6
Period 4 - 4s23d104p6
Period 5 - 5s24d105p6
Period 6 - 6s24f145d106p6
Period 7 - 7s25f146d107p6
1s22s22p63s23p64s23d104p65s24d105p6
6s24f145d106p67s25f146d107p6
Electron Configuration Notation
• Notation used to represent electron
configurations:
– H: 1s1 # of electrons in sublevel/orbital
Energy level
Sub level/orbital
– He: 1s2
– Li: 1s2 2s1
– Be: 1s2 2s2
– B: 1s2 2s22p1
Color the 4 sublevels + make a key
d-1
f-2
…Then write the configuration
of each element!
You Try!
•
•
•
•
•
Write the electron configurations for the following:
• Ca:
C:
22s22p63s23p64s2
2
2
2
–
(20)1s
– (6): 1s 2s 2p
• Ir:
F:
22s22p63s23p64s2
2
2
5
–
(77)1s
– (9): 1s 2s 2p
104p65s24d105p66s2
3d
Ne:
4f145d7
– (10): 1s22s22p6
• Cm:
Na:
– (96)1s22s22p63s23p64s2
– (11)1s22s22p63s1
3d104p65s24d105p66s2
P:
4f145d106p67s25f7
– (15)1s22s22p63s23p3
• Abbreviated form: shows preceding noble gas and
the configuration of only the last energy level!
– Mg: 1s2 2s22p63s2
• or [Ne] 3s2
– B: 1s2 2s2 2p1
• or [He] 2s2 2p1
– Si: 1s2 2s2 2p6 3s2 3p2
• or [Ne] 3s2 3p2
– Al: 1s2 2s2 2p6 3s2 3p1
• [Ne] 3s2 3p1
– Xe: 1s2 2s2 2p6 3s23p64s23d104p65s24d105p6
• [Kr] 5s24d105p6
• What happens in the fourth period?
– After 4s2, comes 3d10, then 4p6
– Scandium (#21): 1s2 2s2 2p6 3s2 3p6 4s2
3d1
• or [Ar] 4s2 3d1
– Copper is [Ar] 4s2 3d9
– Bromine is [Ar] 4s2 3d10 4p5
• What happens in the sixth period?
– After 6s2, comes 4f14, then 5d10, then 6p6
– Tungsten (W) is [Xe] 6s24f145d4
• Why is 3d on the
4th row after 4s?
• ______ energy than 4s
and ______energy than 5p.
• Why is 4f on the 6th row after 6s?
• ______ energy than 6s
and ______energy than 5d
• Aufbau principle: lowest energy orbitals
are filled first!
• Sublevel order =
• 1s,2s,2p,3s,3p,4s,4p,5s,4d,5p,6s,4f,5d,6p
Orbital Notation Rules
1) Aufbau principle: electrons enter orbitals
of lowest energy first.
2) Pauli exclusion principle: an atomic orbital
may describe at most 2 electrons.
3) Hund’s rule: one electron enters each
orbital until ALL orbitals contain 1 electron
with parallel spins.
Light and Atomic Spectra
13-3
(only pg. 372-375)
PLAY: electromagnetic spectrum song:
<iframe width="420" height="315"
src="https://www.youtube.com/embed/bjOGNVH3D4Y"
frameborder="0" allowfullscreen></iframe>
Electromagnetic Spectrum
• Energy in the form of electromagnetic
radiation (radiant energy) travels in waves
• Waves transfer the energy from one place
to another
• Ex: radio waves, TV, microwave, visible
light, x-rays, gamma rays, infrared, UV
• All forms of radiant energy are part of the
electromagnetic spectrum
Low energy
High Energy
Wavelength + Frequency
• Two main properties of electromagnetic
waves:
1) Frequency
2) Wavelength
• Wavelength is the distance between two
corresponding peaks or troughs.
• Frequency is the number of wave cycles per
second.
• Wavelength is inversely proportional to
frequency
• wavelength
frequency
Wave length
Frequency
Wave length
Frequency
• Higher frequency waves (short wavelength)
have high energy
– Ex: gamma rays, x-rays, ultraviolet rays
– Ex: Violet light in visible spectrum
• Low frequency waves (long wavelength)
have low energy
– Ex: radio waves, microwaves, infrared (heat)
waves
– Ex: Red light in visible spectrum
Light and Atomic Spectra
• Electrons absorb energy and move to
higher energy states/levels
• Electrons give off that energy in the form
of light when they fall back down to lower
energy states, or ground state.
• ALL electromagnetic waves travel at the
speed of light in a vacuum – 300 million
meters per second or (3.0 x 108 m/s)
•When atoms are
energized by an electric
current they emit light.
•When this light is
passed through a prism
they produce an
emission spectrum.
•Each element has its
own unique atomic
emission spectrum
fingerprint
http://phys.educ.ksu.edu/vqm/html/emission.html