The Periodic Law
History of the Periodic Table

By 1860, more than 60 elements had
been discovered

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Masses were not accurately known
In 1860, the first International Congress
of Chemists met in Germany
Stanislao Cannizzaro presented a
convincing method for accurately measuring
the relative masses of atoms
 Chemists were able to agree on standard
masses


Dmitri Mendeleev was writing a
chemistry text
He hoped to organize the elements based
on their properties
 He placed the individual elements and their
properties on cards
 He noticed that when the elements were
arranged according to mass, certain
properties appeared at regular intervals


Mendeleev created a table in which
elements with similar properties were
grouped together (1869)
He reversed Te & I based on their
properties
 He left several empty spaces for
undiscovered elements
 He successfully predicted the discovery and
properties of Sc, Ga, and Ge


In 1911, Henry Mosely was working with
Ernest Rutherford
They were examining the spectra of 38
different elements
 This was related to the number of protons,
he noticed that the properties fit this pattern
better than mass
 This explained why Te & I needed to be
reversed

Modern periodic law states that the
physical and chemical properties of the
elements are periodic functions of their
atomic numbers
 The periodic table is an arrangement of
the elements in order of their atomic
numbers so that elements with similar
properties fall in the same column, or
group

William Ramsay discovered argon in
1894
 Helium had been discovered in the sun’s
emission spectrum in 1868
 To accommodate these elements,
Ramsay had to create the column for the
Noble gases
 In 1898, he also discovered krypton and
xenon
 Radon was discovered in 1900 by
Friedrich Dorn

The lanthanides were recognized as a
group of similar elements in the early
1900’s
 The actinides were also identified and to
save space, they are placed with the
lanthanides down below the main body
of the table

Electron Configuration and the
Periodic Table
Generally the electron configuration of
an atom’s highest occupied energy level
governs the atom’s chemical properties
 Vertical columns are called groups or
families and share similar physical
properties
 There are 7 horizontal periods in the
table


There are group names for many individual
families
 The d block elements are known as the
transition metals

Typical metallic properties

The p block element properties vary
greatly

The metals are usually harder and denser
than s block metals, but softer and less
dense than d block metals
The metalloids (semi metals) have
properties of both metals and nonmetals
 The actinides are all radioactive


Group 17 is known as the halogen group


Most reactive nonmetals
React with metals to form salts
Radiant Energy
Much of our understanding of how
electrons behave comes from studies of
how light interacts with matter
 Until the 1800’s scientists believed that
light was a beam of energy moving
through space in the form of waves
 In the 1900’s they found that light also
behaved like a stream of tiny, fastmoving particles


Light travels in electromagnetic waves


form of electromagnetic radiation
An electromagnetic wave consists of
electric and magnetic fields oscillating at
right angles to each other and the
direction of the wave

All waves, whether they are water
waves or em waves, can be described
in terms of four characteristics
amplitude - height of a wave as measured
from the origin to its crest or peak
 wavelength - (l) distance between
successive crests of the wave



visible light is 400-750 nm (can see these w/
eye)
frequency - (n) number of time the wave
cycles up and down in 1 second

expressed as cycles per second (hertz - Hz)


1/s or s-1 (FM radio is in MHz)
speed - (c)in a vacuum is 3.00 x 108 m/s
The relationship between wavelength and frequency is:
l=c/n
n=c/l

The visible spectrum is an example of a
continuous spectrum
one color fades into the next color
 violet has the shortest wavelength, highest
frequency
 red has the longest wavelength, lowest
frequency


visible light is only a small part of the
total electromagnetic spectrum

the rest is invisible to the human eye
Vis
UV
Xray
Quantum Theory
At the beginning of the 20th century, the
wave model of light was universally
accepted
 several observations brought this
acceptance into question

why do hot objects emit light of different
colors as they heat up (red>yellow>white)
 why do elements burn with different colors


In 1900, Max Planck was able to predict
accurately how a spectrum changes w/
T
to do this he proposed that there is a
fundamental restriction on the amounts of
energy that an object emits or absorbs,
and he called each of these pieces of
energy a quantum
 He related the frequency of wavelength
with its energy

E = hn
 h is Planck’s constant =
6.6262 x 10-34 Js
(Joule-seconds)

Using Planck’s theory, scientists can
determine the temperature of far off
objects (ex: stars) by observing their l
 Energy is absorbed or emitted in quanta

hard to imagine in our world (ex: car)
 each quantum is very small, so energy
seems continuous to us
 these quanta can be significant on the
atomic level


Planck’s discovery did not attract much
early attention

Albert Einstein saw a way to explain a
puzzling phenomenon called the
photoelectric effect

For each metal, a minimum frequency of
light is needed to release electrons


regardless of the light intensity
Einstein said that when a photon (quanta)
strikes a surface, it transfers it’s energy to
an electron
electrons take all the energy or none
 higher frequency light has higher energy
photons that can release electrons to flow
 low frequency (low energy) photons cannot
release electrons

This is why x-rays are damaging
and radio waves are not
Einstein won the Nobel Prize in
1921 for explaining this in 1907

In 1923 Arthur Compton convincingly
proved that light consists of tiny
particles (photons)
He demonstrated that a photon could
collide with an electron
 Light therefore has a dual nature


Particles & Waves
Another Look at the Atom

A spectrum that contains only certain
colors, or wavelengths, is called a line
spectrum
every element has a unique line spectrum
when it is heated or electricity is passed
through it
 also called an emission spectrum


Incandescent light gives the complete
spectrum
Emission Spectrum
Absorption Spectrum
Sun
Spectra of different star types
The hottest stars are on top, coolest on the bottom
Our sun is toward the lower middle

Neils Bohr was able to explain why
elements give line spectra
In 1911 he attended a lecture where Ernest
Rutherford was explaining his planetary
model of the atom
 Bohr realized that Planck’s idea of
quantization could be applied to this model
to explain the line spectra

he started with hydrogen
 each electron was allowed to have only certain
orbits corresponding to different energy levels
 he gave each orbit a quantum number (n)
 the lowest orbit was the ground state (n=1)

When an electron absorbs an appropriate
amount of energy, it jumps to a higher
orbit, or excited state (n=2,n=3, n=4, etc.)
 radiation is then emitted as the electron
falls back to the ground state
 Bohr was able to use this model and
Planck’s equation (E=hn) to calculate the
frequencies observed in the line spectrum
of hydrogen
 his model worked well for hydrogen, but
could not explain the spectra of atoms w/
more than one electron, except in a rather
approximate way
 it represented an important initial step in


Until 1900, scientists believed that there
was a clear distinction between matter
& energy


Planck, Einstein, & Bohr showed that
waves had particle properties
In 1924 Louis de Broglie wondered if
matter had waves (matter waves)
eventually proven correct, won the Nobel
Prize, now used in electron microscopes
 we are not aware of this, because they are
so small

a golf ball at 40m/s would have a wavelength of
3X10-34 m
 become significant on the atomic level

Electron micrograph of scratch & sniff paper w/ tiny glass
capsules containing the scent
Louis deBroglie
Werner Heisenberg
-In 1927, Werner Heisenberg proposed the
uncertainty principle
-the position and the momentum of a moving object
cannot simultaneously be measured and known
exactly
-it is not appropriate to think in terms of electrons
moving in well defined orbits, because there is no
way to test this idea
A New Approach to the Atom
The quantum-mechanical model
explains the properties of atoms by
treating the electron as a wave that has
quantized energy
 It is impossible to describe the exact
positions of electrons or how they are
moving
 Describes the probability that electrons
will be found in certain locations around
the nucleus

The probability of finding an electron in
various locations around the nucleus
can be pictured in terms of a blurry
cloud of negative energy
 Cloud is referred to as electron density
 The probability of finding electrons in
certain regions of an atom is described
by orbitals
 An atomic orbital is a region around the
nucleus of an atom where an electron
with a given energy is likely to be found

Orbitals have characteristic shapes,
sizes, and energies
 They do not describe how the electron
actually moves
 Rather than drawing electron clouds to
represent orbitals, it is more convenient
to merely draw the surface within which
an electron is found 90% of the time


Different kinds of orbitals have different
shapes

s is spherical, p is dumbell shaped, d & f
are more complex

The main energy level or principal
energy level in an atom is designated by
the quantum number n


the higher the n, the greater the energy
level
Each principal energy level is divided
into one or more sublevels
the number of sublevels is equal to the
principal quantum number
 the first energy level has just the s sublevel
 the second energy level has an s & p
sublevel


Each sublevel has a specific number of
The higher
the n, the
larger the
orbital
The three p orbitals combine
into the sublevel at the bottom
The 5 d sublevels
*Each orbital can contain 2 electrons
-these will spin either clockwise or
counterclockwise
-to exist in the same orbital they must be
spinning in opposite directions
(opposite magnetic fields)
-Called the Pauli exclusion principle after
Wolfgang Pauli (1900-1958) Austrian
Electron Configurations
The distribution of electrons among the
orbitals of an atom is called the electron
configuration
 Helps chemists to understand chemical
behavior


The Aufbau principle (German for
building up)


Pauli exclusion principle


Electrons are added one at a time to the
lowest energy orbitals available until all the
electrons of an atom have been accounted
for
an orbital can hold a maximum of 2
electrons, must spin in opposite directions
Hund’s rule

electrons occupy equal energy orbitals so
that a maximum number of unpaired
electrons results
Open arrangement of
atomic orbitals
Arrangement for Carbon
Orbital diagrams & configurations for several elements
Notice the sum of the superscripts in the configuration is
equal to the number of electrons

The electron configurations represent
the ground state of the electrons

heating or electric current will cause
electrons to “jump” to higher levels


when they fall, they emit specific amounts of
energy, creating the line spectra discussed
earlier
Cr & Cu are exceptions to the Aufbau
Principle due to the unique stability of a
full or half full sublevel