Chapter 5 Electrons in Atoms
Wave Nature of Light
A form of energy, electromagnetic radiation exhibits wavelike properties. Visible light is
an example of electromagnetic radiation. A wave has certain characteristics.
Wavelength = distance between crests or troughs, symbol λ
Amplitude = height of wave
Frequency, ν = number of waves that pass a given point per second. The units of waves
per second are hertz (Hz) - SI units. Frequency is usually expressed in
“waves per second” or 1/s or s-1
e.g. 300 Hz = 300 waves per second = 300 s-1 (note the word cycle or wave is omitted)
or
300 cycles per second
All electromagnetic waves travel at a speed of light, which is 3.00 x 108 m/s, given the
symbol c.
Useful formula
c=λν
speed of = wavelength * frequency
light
Electromagnetic spectrum p.120
Practice problems 1-4, page 121, page 147 #68, 67
Electromagnetic Spectrum (cont)
Which types of waves or rays have the highest energy?
If the wavelength of microwave radiation is 3 x 10-2m, will the frequency of this radiation
be in the high or low end of the range in the electromagnetic spectrum?
Therefore, we can say that


1

wavelength is inversely proportional to frequency
Particle Nature of Light
Failings of the wave model of light – does not explain why:1) Heated objects emit light of certain frequencies only.
2) When colored light of a specific frequency shines on some metals then electrons
are emitted. This is known as the photoelectric effect, which can be used in
calculators. Here the incident light is converted by photoelectric cells to electrical
energy.
Einstein proposed that electromagnetic radiation exhibits both wave and particle
properties. Light can exist as waves or a stream of tiny particles called photons. Photons
can be thought of as packets of energy; so one photon carries a quantum of energy.
Einstein proposed the following relationship between energy and frequency.
Ephoton =
h
h= 6.626 x 10-34 J.s, Planck’s constant.
p 124 Practice problem 5 and 6.
Atomic Emission Spectra
When atoms are excited by for example heat or electricity, then the atoms absorb energy,
which is later emitted as light. If this light is then passed through a prism, it is split up
into light of different frequencies and produces an ATOMIC EMISSION SPECTRUM – which
is a series of individual lines. Each element has its own atomic emission spectrum.
Balmer Series of spectral lines (visible) transition to the n=2 level
An H electron is normally located in the ground state in the n=1 level. When excited the
electron can drop from higher energy levels to a lower energy level (n=2 in diagram).
When electrons do this, energy is emitted in the form of light. Each transition or drop has
a certain energy associated with it and produces light of a certain frequency from E=h.
When electrons drop back to the n=1 level then E and  are high and light emitted is in
the ultra violet region of the spectrum (lines known as the Lyman series). Transitions
down the n=3 level give spectral lines in the infrared region (Paschen series).
Hydrogen’s Atomic Emission Spectrum
Bohr Model of the Atom
Bohr believed that electrons traveled in circular orbits around the nucleus. The hydrogen
atomic emission spectra is composed of discrete lines that correspond to a certain
frequency of light emitted. The Bohr model correctly predicted the frequencies of the
spectral lines for hydrogen but failed to explain the spectrum of other elements.
The Quantum Mechanical Model of the Atom.
Electrons as waves
Loius de Broglie (1892-1987) predicted that all moving particles have wave
characteristics. He derived an equation for the wavelength () of a particle of mass (m)
moving at velocity (v).
=
h
mv
h= Planck’s constant.
The Heisenberg Uncertainty Principle:
states that it is impossible to know the exact position and velocity of an electron.
Schrődinger studied the wave properties of an electron. Electrons are located in atomic
orbitals, which surround the nucleus of the atom. The Quantum Mechanical Model
assigns quantum numbers to the atomic orbitals:-
The principal quantum number, n, describes the relative size and energy of atomic
orbitals – so the principal quantum number describes the energy level and distance from
the nucleus – n can have values 1, 2, 3 … n.
Electronic Arrangement and Quantum Numbers
The electrons fill up the energy levels starting at the lowest energy level. The formula for
finding the maximum number of electrons in a particular energy level is 2n2, where n =
no. of the energy level.
No. of Energy Level
Maximum no.of e-s.
Pauli Exclusion Principle
This states that no two electrons in any atom have the same amount of energy associated
with it and therefore cannot follow the same path.
Therefore considering the first energy level, n= 1 ( n is the first quantum number),
contains 2 electrons (maximum) these electrons have different spins :- one clockwise and
one anti-clockwise. Different spins have different amounts of energy associated with
them and therefore follow a different path. We denote this as ↑↓.
In each energy level, there can be several sub-levels, 2nd quantum number, l, with levels
s, p, d, f. The s level can take a maximum of two electrons, the p level – 6 electrons, the d
level - 10 electrons and the f energy level -14 electrons.
The letter describing the sub-level is preceded by the value of n for the energy level e.g.
3d, 4f, 2p.
Each sublevel contains orbitals (space occupied by a pair of electrons). The third
quantum number, ml (magnetic quantum number) is used to describe each orbital within
in a sublevel. There are 3 orbitals for the p sublevel, 5 orbitals for the d sublevel, and 7
orbitals for the f sublevel.
Finally, the fourth quantum number ms describes direction of electron spin :- clockwise
or anticlockwise.
Note: No two electrons in an atom can have the same set of quantum numbers.
Relative Stability of Electron Configurations
General Rule:
Atoms with full outer levels are particularly stable. (that is less reactive)
Usually the outer level is full when it contains 8 electrons (exception He: the outer level
is full with 2 electrons).
Note: the helium atom has a full outer level of 2 electrons and He is very unreactive and
stable.
The 8 electrons are made up of 2 electrons in the outer s sublevel and 6 in the outer p
sublevel. These 8 outer level electrons are responsible for the lack of reactivity of the
atom. This is known as the octet rule.
Other Stable Configurations:
An atom with a half filled or completely filled sublevels are also stable.
Example: Cr (Z=24) has a configuration of [Ar] 4s13d5. Here the d orbital is half filled.
And, Cu (Z=29) has a configuration of [Ar]4s13d10. Here the d orbital is completely
filled. Note, we would have expected a configuration of [Ar]4s23d9, but this is less stable.
Order of Stability:
Full outer level > Full sublevel > Half full sublevel > No special arrangement
Properties of Metals and Non-Metals
Metals
Non Metals
Hard, shiny solids
Gases or brittle solids
If solid, surfaces are dull
Conductors of heat and electricity
Insulators, non-conductors of heat and
electricity
Usually high melting and boiling points
Usually low melting and boiling points
Have 3 or fewer electrons in the outer
level
Have 5 or more electrons in the outer
level
On left hand side of Periodic Table –
most elements are metals
On right hand side of Periodic Table
Metalloids have properties of metals and non-metals e.g. B - boron, Si – silicon, Ge –
germanium, Sb – antimony, Te – tellurium,, Po – polonium, and At - astatine