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CH1120 Electronic Structure
of the Atom
Quantum Mechanics The Behavior of the Very Small
Electrons are incredibly small.
Electron behavior determines much of the behavior of atoms.
Directly observing electrons in the atom is impossible, the electron is so small that
observing it changes its behavior.
The Electromagnetic Spectrum
Light is a form of electromagnetic radiation composed of perpendicular oscillating waves, one for the electric field and one for the magnetic field.
Characterizing Waves
The frequency (ν) is the number of waves that pass a point in a given period of time.
The number of waves = number of cycles units are hertz (Hz) or cycles/s = s−1
1 Hz = 1 s−1 or 1/s
The total energy is proportional to the amplitude of the waves and the frequency.
The larger the amplitude, the more force it has. The more frequently the waves strike, the more total force .
low frequency long wavelength
higher frequency
higher frequency short wavelength
low amplitude
higher amplitude
higher amplitude
Characterizing Waves
Increasing energy
The field of quantum mechanics began with the studies of physicists in the early the 20th century.
Max Planck (1918) Albert Einstein (1921)
Neils Bohr (1922)Arthur Compton (1927)Louis de Broglie (1929)
Werner Heisenberg (1932)P. A. M. Dirac (1933)
Erwin Schrödinger (1933)
The Beginnings of Quantum Mechanics
The Relationship Between Wavelength and Frequency
The shorter the wavelength, the more frequently waves pass, and the higher the frequency.
Wavelength and frequency of electromagnetic waves are inversely proportional.
Because the speed of light is constant (3.00 x 108 m/sec), if we know wavelength we can find the frequency, and
vice versa.
ν =cλ
ν ∝ 1 λ
The proportionality constant is c, the speed
of light.
The Relationship Between Wavelength and Frequency
c = 3 x 108 m/s
1. Calculate the frequency (in MHz) of a radio signal with a wavelength of 2.98 m.
ν = 3.00 x 108 m/s2.98 m
= 1.01 x 108 s-1
= 1.01 x 108 Hz
1.00 MHz106 Hz
1.01 x 108 Hz x = 101 MHz
2. Calculate the wavelength of red light with a frequency of 4.62 x 1014 s−1 .
=
3.00 x 108 m/s4.62 x 1014 /s = 6.49 x 10-7 m
6.49 x 10-7 m x 1 nm1 x 10-9 m = 6.49 x 102 nm
RedOrangeYellowGreenBlueViolet
The Electromagnetic Spectrum
The Electromagnetic Spectrum
Energy Increases
Shorter wavelengths of electromagnetic radiation have more energy than longer wavelengths:
Radiowaves have the lowest energy. Gamma rays have the highest energy.
3. Order the following types of electromagnetic radiation:microwaves, gamma rays, green light, red light, ultraviolet light.
By wavelength (short to long)
By frequency (low to high)
By energy (least to most)
gamma < UV < green < red < microwaves
microwaves < red < green < UV < gamma
microwaves < red < green < UV < gamma
White Light Produces a “Continuous” Spectrum
Atomic Spectra
When atoms or molecules absorb energy, that energy is often released as light energy.
When that emitted light is passed through a prism, a pattern of particular wavelengths of
light is seen that is unique to that type of atom or molecule – the pattern is called an
emission spectrum.
non-continuous can be used to identify the material
Emission Spectrum
Red line 𝛌 = 656.3 nm
Green line 𝛌 = 486 nm
Blue line 𝛌 = 434 nm
Violet line 𝛌 = 410.1 nm
Identifying Elements with Flame Tests
Na K Li Ba
Rutherford’s Nuclear Model
The atom contains a tiny dense center called the nucleus.
The nucleus is essentially the entire mass of the atom.
The nucleus is positively charged.
The positive charge balances the negative charge of the electrons.
The electrons move around in the empty space of the atom surrounding the nucleus.
Problems with Rutherford’s Nuclear Model of the Atom
Electrons are moving charged particles.
Moving charged particles give off energy.
Therefore electrons should constantly be giving off energy.
The electrons should lose energy, crash into the
nucleus, and the atom should collapse!!
The Bohr Model of the AtomNeils Bohr (1885–1962)
The energy of the atom is quantized.
The amount of energy in the atom is related to the electron’s position in the atom.
Quantized means that the atom could only have very specific amounts of energy.
Bohr correlated these allowed energy levels with allowed radii of electron orbits.
Bohr Model The energy of each Bohr orbit, specified by a quantum number, n = 1, 2, 3 is fixed, or quantized.
It is impossible for an electron to exist between orbits in the Bohr model.
Bohr orbits are like steps of a ladder, each at a specific distance from the nucleus.
Bohr’s Model
The electrons travel in orbits that are at a fixed distance from the nucleus (stationary states).
The energy of the electron is proportional to the distance the orbit was from the nucleus.
Electrons emit and absorb radiation when they move between orbits.
Emitted radiation is a photon of light. The distance between the orbits determines the
energy of the photon of light produced.
Bohr’s Model
n=1
n=2
n=3 n=4 n=5
When an atom
absorbs energy, an electron is excited to a
higher-energy orbit.
The electron relaxes to a
lower energy level, emitting a photon of light.
Energy is absorbed !!
Energy is emitted !!
Excitation and Emission
Emission vs. Absorption Spectra
Spectra of Mercury
Emission Spectrum
Red line 𝛌 = 656.3 nm
Green line 𝛌 = 486 nm
Blue line 𝛌 = 434 nm
Violet line 𝛌 = 410.1 nm
1
The energy of photons is related to wavelength by another equation:
ℏ = 6.626 x 10-34 Jᐧs
Bohr’s equations related energy of photons to
basic energy shells (n):
n=1
n=2
n=3 n=4 n=5 n=6
Red line 𝛌 = 656.3 nm
Green line 𝛌 = 486 nm
Blue line 𝛌 = 434 nm
Violet line 𝛌 = 410.1 nm
Balmer Series
The Hydrogen Spectrum
Bohr Model The Bohr model also showed that each “principal” energy level could hold a maximum number of electrons.
This explains the increasing length of the “rows” of the
periodic table.
Bohr Model ---> Increasing Length of Periods
Bohr Model ---> Increasing Length of Periods
28
1832
Behavior of main-group elements can be explained in terms of “outer” electrons.
I II III IV V VI VII VIII
adding electrons
The Bohr Model: Atoms with Orbits
The great success of the Bohr model of the atom was thatit predicted the lines of the hydrogen emission spectrum.
However, it failed to predict the emission spectra of other elements that contained more than one electron.
For this and other reasons, the Bohr model was replaced with a more sophisticated model called the quantum-mechanical or
wave-mechanical model.
In the quantum-mechanical model, electrons do not behave like particles flying through space.
We cannot, in general, describe their exact paths.
An orbital is a probability map (a mathematical
model)that shows where the electron is likely to be found when the atom is probed. It does not represent an exact path
that an electron takes.
Electrons reside in principal energy levels which are subdivided into energy sublevels.
The principal energy levels (1-7) can “theoretically” contain 2,8,18,32,50,72, and 98 electrons each.
The Quantum-Mechanical Model: Atoms with “Orbitals”
The number of subshells in a given principal shell is equal to the value of n.
Principal Energy Levels are Divided Into Subshells
Subshells Hold Different Numbers of Electrons
Principal quantum number
Electrons in primary
shell
Electrons in subshells}
Electrons in subshells}
Electrons in subshells}
Electrons in subshells}
n=4
n=3
n=2
n=1
Primary energy shells
Principal Energy Levels
f d p s
d p s
p s
s
n=4
n=3
n=2
n=1
Primary energy shells
Energy subshells
Energy
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
“Degenerate” Orbital Energies
Energy
1s
2s
2p
3s
3p
3d 4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d 7s
“Real” Orbital Energies
Bohr Model (shells)
n=1
n=2
n=3n=4
Modified Bohr Model (subshells)
1s
2s
2p3s3p4s3d4p4d4f
Increasing Energy
5s lower in energy than 4d
6s lower in energy than 4f
4s lower in energy than 3d
]
Filling the Orbitals with ElectronsEnergy levels and sublevels fill from lowest energy to high.
s → p → d →f Aufbau Principle
Orbitals that are in the same sublevel have the same energy
No more than two electrons per orbital
Pauli Exclusion Principle
When filling orbitals that have the same energy, place one electron in each before completing pairs.
Hund’s Rule
Electron Configuration & the Periodic Table
s
f
dp
1 2 3 4 5 6 7
Electron Configuration & the Periodic Table
s
f
dp
1 2 3 4 5 6 7
Electron Configuration & the Periodic Table
s
f
d
p
1 2 3 4 5 6 7
Shell being filled = (period number-1)
Shell being filled = (period number-2)
Shell being filled = period number
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Electron Configuration & the Periodic Table
s
f
d p
1 2 3 4 5 6 7
Shell being filled = period number
Shell being filled = (period number-1)
Shell being filled = (period number-2)
He K Pd Be Co Si Pt U
1s 4s 4d 2s 3d 3p 5d 5f
4. What is the highest energy sublevel being filled for each of the following atoms ?
Electron Configuration & the Periodic Table
s
f
d
p
1 2 3 4 5 6 7
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