Topic 1B
ELECTRONS
THE ELECTRON

J.J. Thomson, English (1898-1903)—found that when
high voltage was applied to an evacuated tube, a “ray”
he called a cathode ray [since it emanated from the (-)
electrode or cathode when YOU apply a voltage across
it] was produced.



The ray was produced at the (-) electrode
Repelled by the (-) pole of an applied electric field, E
He postulated the ray was a stream of NEGATIVE particles
now called electrons, e-
COULOMBS LAW

He then measured the deflection of beams of e- to
determine the charge-to-mass ratio
e is charge on electron in Coulombs, (C) and m is its
mass.
 The interaction between charged objects is a noncontact force that acts over some distance of
separation. Charge, charge and distance
 e/m =-1.76 x 108 C/g
 e is charge on electron in Coulombs, (C) and m is its
mass.
 Or F=q1q2/r2
 If q1 and q2 have the same sign the force is repulsive,
when opposite the forces is attractive

BOHR MODEL OF THE ATOM
Neils Bohr originally came up with the idea of
the atom containing an dense nucleus being
orbited by negatively charged electrons.
 Bohr’s idea (the planetary model) suggested
that electrons could only travel in fixed orbits or
shells around the nucleus.

ATOMIC EXCITATION AND DE-EXCITATION
BOHR & THE SPECTROSCOPE



Bohr used a spectroscope and found that gaseous elements
emitted electromagnetic radiation when heated
The light that was emitted in quanta (energy packets) gave a
unique pattern of radiation.
It was discovered that the release of radiation was caused by
electrons in the atom absorbing energy and being promoted to
a shell further way from the nucleus (ie orbit 1 to orbit 2)
EMISSION SPECTRA



When an electron falls back down to the ground state, it
releases the energy that it absorbed when it was promoted to
the higher energy shell (the exited state).
This release of energy creates a line spectrum.
In other words, the energy emitted is observed as light - this is
the source of the lines we see in the emission spectra
Remember
fireworks??
LYMAN, BALMER & PASCHEN
The Lyman series (red lines) is not visible, in
the UV range (n>1 to n=1)
 Balmer series is in the visible range (green
lines) n>2 to n=2
 Others series like Paschen, Brackett and Pfund)
lie in the infrared region

EXAMPLES OF EMISSION (LINE) SPECTRA
Hydrogen:
Helium:
Carbon:
•The wavelengths of light that are emitted are characteristic
for the particular elements that are present in the gas.
•Thus the emission (bright line) spectrum is a fingerprint of the
gas (or gases) that is (are) emitting the light.
•This is because the shells are in fixed positions, the
difference in energies between them (and the wavelength of
the line in its spectra) is also fixed.
•Video!!
IONIZATION ENERGY


Tie this to ionization energy – sometimes
and electron may gain enough energy to
completely over the attraction of the
nucleus and can be ejected by the atom.
The energy required to do this is different
for each electron



all electrons are at a different distance
from the nucleus (plus shielding)
This is ionization energy!
Bohr’s finding are consistent with the
observed pattern of ionization energies.
THE HYDROGEN ATOM
ABSORPTION (DARK LINE) SPECTRA
SUMMARY – THREE TYPES OF SPECTRA
TYPES OF SPECTRUM





emission spectrum—the spectrum of bright lines, bands, or continuous radiation that
is provided by a specific emitting substance as it loses energy and returns to its
ground state OR the collection of frequencies of light given off by an "excited”
electron
• absorption spectrum—a graph or display relating how a substance absorbs
electromagnetic radiation as a function of wavelength
• line spectrum--isolate a thin beam by passing through a slit then a prism or a
diffraction grating which sorts into discrete frequencies or lines
• Johann Balmer--worked out a mathematical relationship that accounted for the 3
lines of longest wavelength in the visible emission spectrum of H. (red, green and
blue lines)
• Niels Bohr connected spectra, and the quantum ideas of Einstein and Planck: the
single electron of the hydrogen atom could occupy only certain energy states,
stationary states
ABSORPTION OF LIGHT
In photon absorption, an electron goes to a
higher energy level.
 The photon has to have just the right energy to
be absorbed.
 Its energy has to be equal to the energy
difference between the two levels.
 Photons of any other energy are not absorbed.
 What goes up must come down. Energy
absorbed must eventually be emitted

ELECTRONIC CONFIGURATION
Bohr’s planetary model only gave us part of the
picture.
 The fixed orbits did not explain all of the electrons’
properties.
 Enter Debroglie- his work suggested that the
electron was not just a particle but also wave like.
 Schrodinger further developed the wave idea and
was able to make predictions about where an
electron might be found in an atom.

FROM PLANETARY TO QUANTUM MECHANICAL
MODEL


Coupled with the Heisenberg uncertainty
principle (remember can’t know the position of the electron and momentum
at the same time) Schrodinger’s work led to the
quantum mechanical model – the 3d
probability of any one electron at any point in
time within each of the shells.
These 3D probability maps are known as orbitals.
ARRANGEMENT OF
ELECTRONS IN ATOMS
Electrons in atoms are arranged as
SHELLS (n)
SUBSHELLS (l)
ORBITALS (ml)
ARRANGEMENT OF ELECTRONS
IN ATOMS
Each orbital can be assigned no more than 2
electrons!
This is tied to the existence of a 4th quantum
number, the electron
number, ms.
spin quantum
ELECTRON
SPIN
QUANTUM
NUMBER,
MS
Can be proven experimentally that electron
has a spin. Two spin directions are given by
ms where ms = +1/2 and -1/2.
ELECTRON SPIN QUANTUM
NUMBER
Diamagnetic: NOT attracted to a magnetic field
Paramagnetic: substance is attracted to a
magnetic field. Substance has unpaired electrons.
QUANTUM
NUMBERS
n ---> shell
1, 2, 3, 4, ...
l ---> subshell
0, 1, 2, ... n - 1
ml ---> orbital
-l ... 0 ... +l
ms ---> electron spin
+1/2 and -1/2
The electrons address!
PAULI EXCLUSION
PRINCIPLE
No two electrons in the same
atom can have the same set
of 4 quantum numbers.
That is, each electron in an atom
has a unique address of
quantum numbers.
QUANTUM NUMBERS & ATOMIC ORBITALS





There are 4 quantum numbers that describe the “address” and “spin” of an
electron.
The value of n limits the possible values of ℓ, which in turn limit the values of mℓ.
• n—principal energy level--1 to infinity. Determines the total energy of the electron.
 It indicates the most probable [within 90%] distance of the electron from the
nucleus.
 It’s a measure of the orbital size or diameter (cloud size). 2n2 electrons may be
assigned to a shell. The larger the value of n the larger the cloud
 Energy levels closer to the nucleus have lower energy. As n increases, the orbital
 becomes larger and the electron spends more time farther from the nucleus. An
 increase in n also means that the electron has a higher energy and is therefore
less
 tightly bound to the nucleus.
It’s simply the Energy level that electron is in.
If it’s a 3s electron, n = 3, if it’s a 4d electron, n = 4, etc.
ANGULAR MOMENTUM






ℓ--angular momentum--0,1,2,....(n-1) electrons w/in
shell may be grouped into subshells [or
sublevels, same thing!], each characterized by its
certain wave shape.
Each ℓ is a different orbital shape or orbital type.
- n limits the values of ℓ to no larger than n-1. Thus, the
number of possibilities for ℓ is equal to n.
(English translation: 3 sublevels for 3rd E level, 4 for 4th
E level, etc.)
- s,p,d,f sublevels - 0,1,2,3 ℓ-values respectively (So,
what do spdf stand for? Sharp, principle, diffuse,
fundamental--early days of atomic spectroscopy)
You can
keep going
from ‘g’
MAGNETIC QUANTUM NUMBER






mℓ--magnetic—assign the “blanks” in orbital notation with zero on
the middle blank and then -ℓ
that means that the range of orbital's is from +ℓ to - ℓ Designates the
orbital (room) where the electron can be found.
• Gives the DIRECTION IN SPACE that the orbital takes.
• specifies to which orbital within a subshell the electron is assigned.
Orbitals in a given subshell differ only in their orientation in space,
not in their shape.
• Values of ml : from –l, ….0, ….+l.
• The middle orbital of a subshell has a value of 0. Orbitals to the left
of the middle orbital have negative numbers; to the right, they have +
numbers.
How to write quantum numbers
WRITING ATOMIC ELECTRON
CONFIGURATIONS
Two ways of writing
configs. One is
called the spdf
notation.
spdf notation
for H, atomic number = 1
1
1s
value of n
no. of
electrons
value of l
WRITING ATOMIC ELECTRON
CONFIGURATIONS
Two ways of
writing configs.
Other is called
the orbital box
notation.
ORBITAL BOX NOTATION
for He, atomic number = 2
Arrows
2
depict
electron
spin
1s
1s
One electron has n = 1, l = 0, ml = 0, ms = + 1/2
Other electron has n = 1, l = 0, ml = 0, ms = - 1/2
See “Toolbox” for Electron Configuration tool.