1 electrons in atoms. 2 have you even wondered why different atoms absorb and emit light of...
TRANSCRIPT
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Electrons in Atoms
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Have you even wondered why different atoms absorb and emit light of different colors?
The transition of electrons within sublevels releases an amount of energy. If this energy corresponds to the visible section of the spectrum, then we observe colors.
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Gamma rays is the most energetic radiation and shortest wavelength. Gamma rays are produced by the sun, by the stars, by some unstable atomic nuclei on earth. Human exposure to gamma rays are dangerous because of the high energy, they can damage biological molecules.
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X-rays can also damage biological molecules, but it requires excessive exposure to that radiation.
Ultraviolet light is an important component of sunlight. It is not as energetic as gamma or X-rays, but still carries enough energy to damage biological molecules. Excessive exposure to ultraviolet light increases the risk of skin cancer and cataracts.
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Infrared light – the heat that you feel when you place your hand near a hot object is infrared light. All hot objects including the human body emit infrared light. It is invisible to our eyes, but infrared sensors can detect the infrared light and these sensors are often used in night vision to “see” in the dark.
Visible light the only light detected by the human eye.
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Microwave light – lower energy but it is absorbed by water and therefore heat substances that contain water. For this reason substances that contain water, such as food, are warmed in a microwave oven, but substances that do not contain water such as a plate are not.
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The longest wavelengths, are used to transmit signals responsible for AM and FM radios, cellular phones, television and other forms of communication.
Radiowaves
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Parts of a wave
l
Wavelength
Wavelength = distance between two consecutive crests. Measure in meters, cm, nmRecall 1 m = 1x109 nmFrequency = number of cycles in one secondMeasured in hertz 1 hertz = 1 /second
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Frequency = n
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• The longer the wavelength (λ), the shorter the frequency (ν).
• The relationship is expressed as:
• c = ν λ
• where c = speed of light (constant) = 3.0x108 m/s
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What is the wavelength of light with a frequency 5.89 x 105 Hz?
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What is the frequency of blue light with a wavelength of 484 nm?
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The Hydrogen Line Emission Spectrum
The lowest energy state of an atom is its ground state.
A state in which an atom has the highest potential energy than it has in its ground state is an excited state.
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The Hydrogen Line Emission Spectrum
Absorption and Emission Spectra
•http://www.flinnsci.com/atomicspectrum
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Energy is Quantized Planck found DE came in “chunks”
with size hn DE = hn and h is Planck’s constant h = 6.626 x 10-34 J s these packets of hn are called
quantum See Planck’s ideas
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• A photon is a particle of electromagnetic radiation with zero mass and carrying a quantum of energy.
• The energy of a particular photon depends on the frequency of the radiation:
Ephoton = h ν
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Niels Bohr Developed the planetarium model of the atom. He said the atom was like a solar system where
electrons rotate around the nucleus like planets around the sun.
According to the model, the electron can circle the nucleus only in allowed paths or orbits.
The energy of the electron is higher when the electron is in orbits that are successively farther from the nucleus
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The Bohr Model of the Atom Energy must be added to an atom in
order to move an electron from a lower energy level to a higher energy level. This is absorption.
When an electron fall to a lower energy level, a photon is emitted and the process is called emission.
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The Bohr Model of the Atom
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Ene
rgy
When energy is put into an atom, the absorbed energy allows the electron to reach higher energy levels. The electron will be in an excited state.
n=1
n=5
n=4
n=3
n=2
n=6n=7
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Ene
rgy Only certain energies
are allowed, so in order for the electron to “jump” to a higher energy level, it will have to absorb the energy equal to the energy gap between the energy levels.
n=1
n=5
n=4
n=3
n=2
n=6n=7
ΔE=E4-E1
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Ene
rgy When an excited
electron drops to a lower energy level, the atom emits energy. The emitted energy is what makes the emission spectra.
n=1
n=5
n=4
n=3
n=2
n=6n=7
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Ene
rgy The energy released by
the electron is directly proportional to the frequency (ν) and inversely proportional to the wavelength (λ)of the radiation.
n=1
n=5
n=4
n=3
n=2
n=6n=7
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Ene
rgy This transition will
show a line in the spectra in a different position (wavelength) as the transition n=7 n=1
n=1
n=5
n=4
n=3
n=2
n=6n=7
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The Quantum Mechanical Model A totally new approach.
Schröedinger’s Equation Developed an equation that treated electrons in atoms as
waves. The wave function is a F(x, y, z) Electrons do not travel around the nucleus in neat orbits,
like Bohr postulated. Instead, the exist in certain regions called orbitals. Orbital is a three dimensional region around the nucleus
that indicated the probable location of an electron.
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The Heisenberg Uncertainty Principle
It is impossible to determine simultaneously both the position and the velocity of an electron or any other particle.
Both, the Heisenberg Uncertainty Principle and the Schrödinger wave equation laid the foundation for modern quantum theory.
Heisenberg Principle
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Quantum theory
Quantum numbers and orbitals
• The solutions to the Schröndinger wave equation are called quantum numbers and they describe the probability of finding the electron around the nucleus.
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S orbital
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P orbitals
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D orbitals
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F orbitals
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F orbitals
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•Electron spin quantum number (s)
•Can have 2 values.
•either +1/2 or -1/2
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Energy level 1
Energy level 2
Energy level 3
Energy level 4 4 sublevels
3 sublevels
2 sublevels
1 sublevel
4f (14 electrons)
3p (6 electrons)
3s (2 electrons)
2p (6 electrons)
2s (2 electrons)
1s (2 electrons)
4s (2 electrons)
4p (6 electrons)
4d (10 electrons)
3d (10 electrons)
2(n)2
2(1)2= 2
2(2)2= 8
2(3)2= 18
2(4)2= 32
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Aufbau Principle Aufbau is German for building up.
Electrons enter orbitals from low energy to high energy.
The order of orbitals based on their energies is:
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p …
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Pauli Exclusion Principle
No two electrons in the same atom can have the same set of quantum numbers.
Even if two electrons are located in the same energy level and same orbital, they must be different in the spin number.
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
H with 1 electron
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
He with 2 electrons
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
Li with 3 electrons
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
Be with 4 electrons
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
B with 5 electrons
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
C with 6 electrons
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
N with 7 electrons
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
O with 8 electrons
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
F with 9 electrons
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
Ne with 10 electrons
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
Na with 11 electrons
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
Mg with 12 electrons
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
Ar with 18 electrons
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
K with 19 electrons
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
Ca with 20 electrons
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
Sc with 21 electrons
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
Mn with 25 electrons
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
Ga with 31 electrons
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Orbital Notation see PPT Valence electrons- the electrons in
the outermost energy levels (not d). Core electrons- the inner electrons. Hund’s Rule- The lowest energy
configuration for an atom is the one have the maximum number of unpaired electrons in the orbital.
C 1s2 2s2 2p2
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• Orbital notation or box notation allows to consider Hund’s rule and the spin of the electron.
For example: 8O 1s2 2s2 2p4
• With orbital notation, the orbitals are represented by boxes (or some authors use circles). That way the s orbital will only be represented by 1 box, p orbitals by 3 boxes, and so on.
8O
1s2 2s2 2p4
• Oxygen is paramagnetic with 2 unpaired electrons
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25Mn1s2 2s2 2p6 3s2 3p6 4s2
3d5
Manganese is paramagnetic with five unpaired electrons.
Paramagnetism is the property of substances of being attracted by an external magnetic field. It is present in all substances with at least 1 unpaired e-.
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30Zn1s2 2s2 2p6 3s2 3p6 4s2
3d10
Zinc is diamagnetic with no unpaired electrons.
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26Fe1s2 2s2 2p6 3s2 3p6 4s2
3d6
Iron is paramagnetic with 4 unpaired electrons
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Details Elements in the same column have
the same electron configuration. Elements in same groups have similar
properties because of similar electron configuration.
Noble gases have filled energy levels. Transition metals are filling the d
orbitals.
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1
H
3
Li11
Na19
K37
Rb55
Cs87
Fr
4
Be12
Mg20
Ca38
Sr56
Ba88
Ra
41
Nb
21
Sc39
Y71
Lu103
Lr
23
V22
Ti40
Zr72
Hf104
Rf
73
Ta105
Db
24
Cr42
Mo74
W106
Sg
25
Mn43
Tc75
Re107
Bh
44
Ru
26
Fe
76
Os108
Hs
27
Co45
Rh77
Ir109
Mt
28
Ni46
Pd78
Pt110
29
Cu47
Ag79
Au111
30
Zn48
Cd80
Hg112
13
Al
83
Bi
51
Sb
33
As
15
P
114
82
Pb
50
Sn
32
Ge
14
Si
10
Ne9
F8
O7
N6
C
86
Rn81
Tl
49
In
31
Ga
5
B
84
Po116
85
At
54
Xe
36
Kr53
I
35
Br52
Te
34
Se
18
Ar
2
He
17
Cl16
S
59
Pr58
Ce61
Pm60
Nd57
La62
Sm63
Eu64
Gd65
Tb66
Dy67
Ho70
Yb69
Tm68
Er91
Pa90
Th93
Np92
U89
Ac94
Pu95
Am96
Cm97
Bk98
Cf99
Es102
No101
Md100
Fm
s block
d block
p block
f block