Transcript 7.1 The Quantum Atom

The Quantum Mechanical
Model of the Atom
Rutherford’s Model of the Atom
In 1911, Ernest Rutherford proposed the nuclear planetary model
of the atom. In this model, most of the atom’s mass was
located in a very small space in the centre of the atom (the
nucleus – positively charged). Around the nucleus, electrons
(negatively charged) orbited like planets orbiting the sun.
Problems with Rutherford’s Model
• In Classical physics, when a charged particle oscillates or
accelerates, it gives off electromagnetic radiation (light energy).
• Since the electron is constantly accelerating (constantly
changing its direction), it should give off light energy.
• But if an electron is giving off energy, it should slow and fall in
towards the nucleus, causing the atom to collapse and to emit
further radiation (calculated to be violet or ultraviolet light),
ending in an “ultraviolet catastrophe”.
• Since atoms do not collapse in a burst of ultraviolet light, the
planetary model had to be incorrect.
Bohr’s Model of the Atom
• Niels Bohr’s Model of the atom indicated that electrons were restricted to
given energy levels or shells around an atom. Any given shell could hold at
maximum, 2n2 electrons, where n = the energy level, a natural number from
1 to 2 to 3 etc. Thus the maximum number of electrons that could be
contained on energy level 4 is 2(4)2 = 32. Bohr’s shells were given number
and letter values: K (n=1), L (n=2), M (n=3), N (n=4) etc.
Bohr’s Model of
the Atom
• Atoms absorb specific
energies by sending electrons
to specific higher energy
levels and atoms emit specific
energies (as light) when
electrons drop to specific
lower energy levels.
Problem With Bohr’s Model: It Only Explained The
Hydrogen Atom
• Bohr’s Model failed to explain atoms besides Hydrogen.
The bright line spectra of other atoms often break apart into
smaller spectral lines in magnetic fields and Bohr’s atom
could not explain this.
Einstein (1900) Suggested That Light Be Thought of as
Energy Particles
• Louis De Broglie
suggested that if
energy could be
thought of as
particles, matter
(electrons) could be
thought of as waves.
Electron “standing
waves” would have
specific wavelengths
depending on their
energy value and
distance from the
nucleus.
The Heisenburg Uncertainty Principle
• Heisenburg demonstrated that it was impossible to know both
the position and speed of an electron since the energy used to
detect this would alter both the position and speed of an
electron, making it impossible to know these.
• This means that it is only possible to determine the probability
of where an electron will be. The shape of the space where an
electron is likely to be is called an orbital (as distinguished from
an orbit which is a circular path, something that can not be
because of the uncertainty principle).
The Heisenburg Uncertainty Principle
Erwin Schrodinger: Treated Electrons as Waves
• Treating electrons as waves and using
complicated wave equations, Schrodinger was
able to produce electron probability
distributions for where an electron was likely to
be located in space around the nucleus.
Where an electron is likely to be (Electron Probability
Distribution) is determined by four quantum numbers
which solve Schrodinger’s equations.
• The four quantum numbers are (n,l,m,s)
• These numbers are put into a wave equation and solve the
equation, producing a three-dimensional representation (orbital)
of where the electron is most likely to be.
• By analogy, we can locate a person on earth by (Country,
Province, City, Street Address) which is like the quantum
numbers (n,l,m,s)
n, the Principal Quantum Number
• The principal quantum
number corresponds to
Bohr’s energy shells.
• n is a natural number
(1,2,3,4,5.etc.)
• n indicates how far from
the nucleus the electron
is likely to be.
• n refers to the major
energy level an electron
will be found in.
• Capital letters (K,L,M,N)
also refer to the major
energy levels.
l , the orbital (azimuthal) number
• The orbital number, l ,
(azimuthal number) refers
to the minor energy level
an electron is in, within its
major energy level.
• The orbital number on a
given major energy level
is a whole number that
goes from 0 to 1 to 2 to a
maximum of n-1.
• For n=1, l = 0. For n=2, l
= 0, 1. For n=3, l = 0, 1,
2.
• On a major energy level,
there are as many minor
energy levels as the value
of n (ex: three minor
energy levels on n=3,
designated by l = 0, 1, 2.
Orbitals
• Orbitals are three-dimensional spaces in which electrons
have a high probability of being found (99% chance).
• A single orbital can hold 0, 1, or 2 electrons.
l , the orbital (azimuthal) number
• Instead of numbers, l is usually denoted by letters.
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l=0 is referred to as s (spherical shaped orbital)
l=1 is referred to as p (dumbbell-shaped orbitals)
l=2 is referred to as d (many-shaped orbitals)
l=3 is referred to as f (many-shaped orbitals)
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Orbital Shapes
l=0 is referred to as s (spherical shaped orbital)
l=1 is referred to as p (dumbbell-shaped orbitals)
l=2 is referred to as d (many-shaped orbitals)
l=3 is referred to as f (many-shaped orbitals)
m, the magnetic quantum number
• m, the magnetic quantum number indicates the orientation of the
orbital(s) in space as well as the number of orbitals on a minor energy
level.
• m is an integer from –l to + l. The values for m depend on the l minor
energy level.
• For l =0 (spherical shape), m=0 (One way to orient a sphere)
• For l =1 (dumbbell shapes), m= -1, 0, +1 (three ways to orient)
• For l =2 (many shapes), m= -2, -1, 0, 1, 2
Overview of Major Energy Levels, Minor Energy
Levels and Orbitals
• The number of minor energy levels on a major energy level is n
• The number of orbitals on a major energy level (pel) is n2
• The number of electrons accommodated on a pel is 2n2
s, the spin number
• s, the spin number, indicates whether an electron has a
clockwise spin or a counterclockwise spin.
• s is a rational fraction of either +1/2 (clockwise) or -1/2
(counterclockwise).
Energy Levels and Orbitals
• The following diagram summarizes the energy levels and
orbitals but does not properly show the relative energies of
minor energy levels,
Energy Levels and Orbitals
• The following diagram summarizes energy levels and
orbitals and shows the relative energies of minor energy
levels,
Electrons in Atoms
• To show electron placement of an atom, determine the number
of electrons from the atomic number (and ion charge for ions).
• Begin filling electrons from lowest to higher energy levels.
• Use Hund’s Rule for multiple orbitals on the same minor energy
level: No orbital can have two electrons until each orbital has at
least one electron.
Electrons in Atoms
• Use the Aufbau Principle to determine the order in which
orbitals fill:
Orbital Diagrams and Electron Configurations
• Two common ways to show electron structure in atoms are
orbital diagrams and electron configurations.
Ground State vs Excited Atoms
• A ground state atom has its electrons in the lowest
possible orbitals and this atom has the lowest energy.
• An excited atom has one or more of its electrons in higher
orbitals so that one or more lower electrons is unpaired or
a lower orbital is empty.
The Orbitals of An Atom Are Overlayed in Space
d orbitals can borrow from lower s orbitals
• When d orbitals are one electron from being half-full or full,
they borrow one electron from the next lowest s orbital to
make a half-full or full d minor energy level. This is
possible because d and s orbitals are very close in energy
value.
• Ex: Cr and Cu
Chemical Families Explained by Quantum
Mechanical Model
• Elements display similar chemical behaviour because they
have the same electron configurations at different energy
levels.
Magnetism Explained by Quantum Mechanical
Model
• When atoms have a series of unpaired d orbitals, the magnetic effects of the
unpaired electrons intensify and lead to the atom itself having a strong
magnetic field.
• Neighboring magnetic atoms together intensify the overall magnetic fields of
each other, creating microscopic regions of atoms called domains. Each
domain has magnetic poles.
• Fe: 1s2 2s22p63s23p64s2 3d23d13d13d13d1
Quantum Mechanical Model and Periodic Chart
• The periodic chart is composed of blocks which reflect the
quantum mechanical model of the atom
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