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Quantum Mechanics. Bohr’s theory established the concept of atomic energy levels but did not thoroughly explain the “wave-like” behavior of the electron.

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Quantum mechanics
Quantum Mechanics

  • Bohr’s theory established the concept of atomic energy levels but did not thoroughly explain the “wave-like” behavior of the electron.

  • Current ideas about atomic structure depend on the principles of quantum mechanics, a theory that applies to subatomic particles such as electrons.


Quantum mechanics1
Quantum Mechanics

  • The first clue in the development of quantum theory came with the discovery of the de Broglie relation.

  • In 1923, Louis de Broglie reasoned that if light exhibits particle aspects, perhaps particles of matter show characteristics of waves.

  • He postulated that a particle with mass m and a velocity v has an associated wavelength.

  • The equation λ= h/mvis called the de Broglie relation.


Quantum mechanics2
Quantum Mechanics

  • If matter has wave properties, why are they not commonly observed?

  • The de Broglie relation shows that a baseball (0.145 kg) moving at about 60 mph (27 m/s) has a wavelength of about 1.7 x 10-34 m.

  • This value is so incredibly small that such waves cannot be detected.


Quantum mechanics3
Quantum Mechanics

  • If matter has wave properties, why are they not commonly observed?

  • Electrons have wavelengths on the order of a few picometers (1 pm = 10-12 m).

  • Under the proper circumstances, the wave character of electrons should be observable.


Quantum mechanics4
Quantum Mechanics

  • If matter has wave properties, why are they not commonly observed?

  • In 1927, it was demonstrated that a beam of electrons, just like X rays, could be diffracted by a crystal.


Quantum mechanics5
Quantum Mechanics

  • Quantum mechanics is the branch of physics that mathematically describes the wave properties of submicroscopic particles.

  • We can no longer think of an electron as having a precise orbit in an atom.

  • To describe such an orbit would require knowing its exact position and velocity.

  • In 1927, Werner Heisenberg showed (from quantum mechanics) that it is impossibleto know both simultaneously.

Heisenberg Uncertainty Principle

It’s impossible to know both the exact position and exact momentum of an e-.


Quantum mechanics6
Quantum Mechanics

  • Heisenberg’s uncertainty principle is a relation that states that the product of the uncertainty in position (Δx) and the uncertainty in momentum (mΔvx) of a particle can be no larger than h/4π.

  • When m is large (for example, a baseball) the uncertainties are small, but for electrons, high uncertainties disallow defining an exact orbit.


Quantum mechanics7
Quantum Mechanics

  • Although we cannot precisely define an electron’s orbit, we can obtain the probability of finding an electron at a given point around the nucleus.

  • Erwin Schrodinger defined this probability in a mathematical expression called a wave function, denotedy(psi).

  • The probability of finding a particle in a region of space is defined byy2.


Plot of y 2 for the lowest energy level of the hydrogen atom
Plot ofy2 for the lowest energy level of the hydrogen atom.



Quantum numbers and atomic orbitals
Quantum Numbers and Atomic Orbitals about the nucleus

  • According to quantum mechanics, each electron is described by four quantum numbers.

  • Principal quantum number (n)

  • Angular momentum quantum number (l)

  • Magnetic quantum number (ml)

  • Spin quantum number (ms)

  • The first three define the wave function for a particular electron. The fourth quantum number refers to the magnetic property of electrons.


Quantum numbers and atomic orbitals1
Quantum Numbers and Atomic Orbitals about the nucleus

  • The principal quantum number(n)represents the “shell number” in which an electron “resides.”

  • The smaller n is, the smaller the orbital.

  • The smaller n is, the lower the energy of the electron.


Quantum numbers and atomic orbitals2

about the nucleus 0 1 2 3 4 5 ...

Code s p d f g h ...

Quantum Numbers and Atomic Orbitals

  • The angular momentum quantum number (l)distinguishes “sub shells” within a given shell that have different shapes.

  • Also known as the Azimuthal quantum number

  • Each main “shell” is subdivided into “sub shells.” Within each shell of quantum number n, there are n sub shells, each with a distinctive shape.

  • l can have any integer value from 0 to (n - 1)

  • The different subshells are denoted by letters.


Quantum numbers and atomic orbitals3
Quantum Numbers and Atomic Orbitals about the nucleus

  • The magnetic quantum number (ml)distinguishes orbitals within a given sub-shell that have different shapes and orientations in space.

  • Each sub shell is subdivided into “orbitals,” each capable of holding a pair of electrons.

  • ml can have any integer value from -l to +l.

  • Each orbital within a given sub shell has the same energy.


Quantum numbers and atomic orbitals4
Quantum Numbers and Atomic Orbitals about the nucleus

  • The spin quantum number (ms)refers to the two possible spin orientations of the electrons residing within a given orbital.

“spin down” or ↓

  • ms = +½ or −½

“spin up” or ↑

Spinning charges generate magnetic fields

Pauli: every e- in an atom must have a unique set of (n, ℓ, mℓ, ms).

  • Max. of 2e- per orbital (with opposite spins: ↑↓).

  • Pauli exclusion principle.


Electron spin
Electron Spin about the nucleus

Experiments showed a spin quantum number, ms

was also needed

  • ms = +½ or −½

“spin down” or ↓

“spin up” or ↑

Spinning charges generate magnetic fields

Pauli: every e- in an atom must have a unique set of (n, ℓ, mℓ, ms).

  • Max. of 2e- per orbital (with opposite spins: ↑↓).

  • Pauli exclusion principle.



The 2 p orbitals
The 2 about the nucleusp orbitals.


The five 3 d orbitals
The five 3 about the nucleusd orbitals.


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