The Evolution of the Atom
1808: Dalton’s model of the atom was the billiard ball model. He
thought the atom was a solid, indivisible sphere. Atoms of each
element were identical in mass and their properties. Atoms of one
element differed from that of another atom. P 866
1897: JJ Thomson’s model of the atom was the plum pudding model.
He discovered the electron and believed that a massive positively
charged substance filled the atom. The electrons were arranged
within this substance. P 867
1898: Ernest Rutherford’s model of the atom was the nuclear model.
He directed α particles at a thin sheet of gold foil. Most of the α
particles passed through the foil but some were deflected at large
angles. These results could only be explained if all the positive charge
of the atom was concentrated in a tiny, massive central core called
the nucleus. P 868
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an atom is 10000 times larger than the nucleus. It is mostly
empty space.
1911: Neils Bohr’s began work on the atom. There was an
inconsistency with the nuclear atom and classical theory. Rutherford’s
model did not account for the lack of emission of radiation as
electrons move about the nucleus (at the rate an electron would lose
energy, it would spiral into the nucleus) and the emission of light at
only certain wavelengths (an accelerated electron should radiate
energy at all wavelengths). P 868
Bohr tried to unite the nuclear mode with Einstein’s quantum theory
of light. He proposed characteristics for the atom that appeared to
contradict classical laws because he realized some phenomena are
unobservable on a macroscopic level and are apparent on an atomic
level.
Light emitted by atoms had been studied. The set of light
wavelengths emitted by an atom is called the atom’s emission
spectrum.
Hydrogen was studied extensively. It produced a unique spectrum of
light. This very precise set of frequencies of EMR extended from the
infrared region well into the UV region of the spectrum. P 870
Bohr thought the emission spectra was too complex to be useful. But
when introduced to a pattern developed by Balmer, the whole model of
the atom became clear.
The Balmer series works for the spectral lines of hydrogen that lie in
the visible range. P 871
Rydberg modified Balmer’s formula to incorporate all possible lines in
the hydrogen spectrum. P 871
Bohr’s Postulates
- Electrons exist in circular orbits. It is the electrostatic force
that holds them in orbit rather than a gravitational force.
- Electrons only exist in allowed orbits. In each orbit there is a
total amount of energy so these orbits are described as energy
levels. This means that the energy of electrons in atoms is
quantized.
- If an electron remains in orbit, it does not radiate energy.
- When electrons jump between orbits, they absorb or emit an
amount of energy that is equal to the difference in the energy
levels.
Bohr said the change in energy of an electron when a photon is
absorbed is equal to the energy of a photon.
hf = E excited − E ground = E f − E i
So, when the electron is returned to the ground state, a photon is
emitted.
Bohr found that for the hydrogen atom, the energy associated with a
particular level was given by:
En = −
13 . 6 eV
n2
Where n is the principle quantum number (energy level) and energy is
negative because energy is being added to pull the electron away from
the nucleus.
Quantum Physics
Bohr’s model of the atom was a first step into the quantum nature of
the atom BUT, it was incomplete. Very close examination of the lines
in the hydrogen spectrum showed that some of the lines were made up
of several fine lines that were very close together. P 877
Another feature that the Bohr atom could not explain was that when a
sample was placed in a magnetic field, spectral lines would split. What
had been one line in the spectrum became two or more lines. This is
known as the Zeeman effect.
Several physicists attempted to modify the Bohr atom without
success. It was De Broglie’s concept of matter waves that paved the
way to the new quantum mechanics or wave mechanics.
When electrons were moving in circular orbits around the nucleus, the
pilot waves must form standing waves otherwise they would be
eliminated by destructive interference. P 878
1925: Schrodinger read de Broglie’s thesis and within weeks, he
developed a very complex mathematical equation to produce detailed
information about matter waves and the atom. The Schrodinger wave
equation forms the foundation of quantum mechanics. When the wave
equations are solved, you obtain wave functions. Wave functions
provide information about the allowed orbits and energy levels of
electrons in the atom.
From this, two more quantum numbers were obtained.
Principle Quantum Number, n – specifies the energy level of the
electron.
Orbital Quantum Number, l – specifies the shape of the orbital.
Values of l are non-negative integers and are less than n.
l (0, 1, 2, 3, …) are assigned the letters s, p, d, f, … p 881
Magnetic Quantum Number, ml – specifies the orientation of the
orbitals when the atom is placed in a magnetic field.
Spin Quantum Number, ms – specifies the orientation of the electron.
It can only be +1/2 or –1/2.
Two questions:
1. If most of the mass of an atom is in the nucleus and the electrons
are in “clouds” that are enormous compared to the nucleus, why is
matter so solid?
2. Why cannot atoms be compressed into much smaller volumes?
1925: Pauli answered these questions which is known as the Pauli
exclusion principle. It states that no two electrons in the same atom
can have the same four quantum numbers. Electron clouds cannot
overlap. P 883