Electrons and
the Periodic Table
Honors Chemistry
The
Periodic
Table
Development of the Periodic Table
1) History of the Periodic Table – By the end of the
1700’s, scientists had identified only 30 elements
(ex. Cu, Ag, Au, H2, N2, O2, C).
2) By the mid 1800’s, about 60 elements had been
identified.
3) Sept 1860 – chemists assembled at the First
International Congress of Chemists in Germany
to settle the controversial issues such as atomic
mass. Standard values set for atomic mass and
improved communication for research.
Johann Dobereiner: 1817
Organized the elements into sets
of three with similar
properties.
He called these groups triads.
The middle element is often
the average of the other two.
Ex)
Cl – 35.5
Br – 79.9
I – 126.9
Ca
Avg
Sr
Ba
Cl + I
 Avg.
2
Triads on the Periodic Table
John Newlands: 1866
• Arranged elements in
order of increasing
atomic mass.
• Noticed repeating
patterns in the elements’
properties every 8th
element.
• Law of Octaves properties of elements
repeated every 8th
element.
• There were 62 known
elements at the time.
Dmitri Mendeleev: 1869
•
•
•
•
Arranged elements in order of
increasing atomic mass.
Similar properties occurred
after periods (horizontal rows)
of varying lengths.
Organized the 1st periodic
table according to increasing
atomic mass and put elements
with similar properties in the
same column.
Periodic – repeating properties
or patterns
Mendeleev’s
1st Periodic
Table
•
•
•
•
•
Mendeleeve Noticed inconsistencies in the
arrangement of increasing atomic mass.
He arranged some elements out of atomic
mass order to keep them together with other
elements with similar properties. (Notice Te
and I)
He also left several blanks in his table.
In 1871, he correctly predicted the existence
and properties of 3 unidentified elements –
Sc, Ga and Ge
These elements were later identified and
matched his predictions.
1st Periodic Law
Properties of the elements repeat
periodically when the elements are
arranged in increasing order by atomic
mass
Mendeleev is known as the
Father of Chemistry
Element 101 (Md) honors
Mendeleev
What’s wrong with
Mendeleev’s PT?
• We know that Mendeleev's periodic table
was underpinned by false reasoning.
– Mendeleev believed, incorrectly, that chemical
properties were determined by atomic weight.
• In 1869 the electron itself had not been
discovered until 1896 - 27 years later.
• It took 44 years for the correct explanation
of the regular patterns in Mendeleev's
periodic table to be found...
Henry Moseley: 1913
1. Fired electrons at atoms, resulting in
emission of x-rays. Each element
emitted x-rays at a unique frequency.
–
He examined the pattern that was best
explained if the positive charge in the
nucleus increased by exactly one unit from
element to element.
2. Element are different from one another
because their atoms have different
number of positive charge (protons).
3. Analyzed data and found that the
elements in the PT fit into patterns better
when arranged in increasing nuclear
charge, which is the Atomic Number.
The Modern
Periodic Table:
When elements
are arranged in
order of
increasing
atomic number,
their physical
and chemical
properties show
a periodic
pattern.
Glenn Seaborg
“Seaborgium” Sg #106
• Born in 1912 in Michigan,
Seaborg proposed reorganizing
the Periodic Table one last time
as a young chemist working on
the Manhattan Atomic Bomb
Project during WWII.
• He suggested pulling the “fblock” elements out to the
bottom of the table.
• He was the principle or codiscoverer of 10 transuranium
elements.
• He was awarded the Noble
prize in 1951 and
died in 1999.
Seaborgium is the exception…
• After some argument between the USA and the rest of the
world, element 106 was named Seaborgium shortly before he
died. This was a matter of some controversy because the
International Union of Pure and Applied Chemistry, IUPAC,
the body that deals with naming in chemistry, had previously
ruled that elements should not be named after living people.
Atomic #
104
105
106
107
IUPAC
Unnilquadeum
Unq
Unnilpentium
Unp
Unnilhexium
Unh
Unnilseptium
Uns
Agreed in
1995
Dubnium
Joliotium
Rutherfordium
Bohrium
(Dubna, Russia)
(Frederic Joliot)
(Earnest Rutherford)
(Neils Bohr)
Rutherfordium
Dubnium
Seaborgium
Bohrium
Agreed in
1996
Parts of the Periodic Table
A. Horizontal Rows – PERIODS
– There are 7 periods in the periodic table
– Elements in a period do NOT have similar
properties.
B. Vertical Columns – GROUPS or FAMILIES
– Labeled 1-18
– IA-VIIIA are the Main-group or representative
elements.
– Elements in a group have similar properties.
– Why?
Family Names
write these on your P.T.
• Transition elements or
• Hydrogen (1)
metals (3-12): d-block
• Alkali metals (1) – most
• Inner transition
reactive metals;
elements or metals (freactivity increases
block)
down the group
– Lanthanides or
• Alkaline earth metals (2)
lanthanide series
• Boron family (13)
– Actinides or actinide
series
• Carbon family (14)
– Transuranium elements
• Nitrogen family (15)
• Oxygen or Chalcogen
family (16)
• Halogens (17)
• Noble gases (18) - inert
1A
1
3A
13
2A
2
4A 5A 6A 7A
14 15 16 17
8A
18
Transition Metals
3
4
5
6
7
8
9
10 11 12
Inner Transition Metals
Lanthanide Series
Actinide Series
1 = Alkali Metals and Hydrogen Group
15= Nitrogen Group
2 = Alkaline Earth Metals
16 = Oxygen or Chalcogen Group
13 = Boron Group
17 = Halogen Group
14 = Carbon Group
18 = Noble Gas Group
Parts of the Periodic Table
The ELECTRON
How are the
electrons
arranged
around the
nucleus?
To Review
• Dalton  Thomson  Rutherford  Bohr
 Quantum Mechanical (Schrödinger) Model
• Bohr – electrons in a particular path have a
fixed energy called energy levels
– Rungs of a ladder
• Quantum Mechanical (Schrödinger) Model
– Electrons better understood as WAVES
– Does not tell where the electrons are located
– Electrons have a certain amount of energy - QUANTIZED
Parts of a Wave
Light as a Wave
Characteristics of a Wave
A. Amplitude: Height of the wave from the baseline.
The higher the wave the greater the intensity.
B. Wavelength: (λ , “lambda”) in nanometers
(1 x 10-9 m). Distance between similar points on 2
consecutive waves.
C. Frequency: (ν , “nu”) The number of waves that
pass a fixed point per unit of time. Measured in
cycles/second (1/s) 1 cycle/second = Hertz (Hz)
Baseline
Wavelength
Crest
Amplitude
Wavelength
Amplitude
Trough
3
D. Electromagnetic Radiation
- a form of energy that exhibits wavelike behavior as
it travels through space
- all forms of EM radiation move at the speed of light
Speed of Light (c)
E. 3.00 x 108 m/s or 186,000 miles/sec.
The relationship between wavelength and
frequency can be shown with the following
equation:
c=λν
This is an indirect relationship.
If λ  then ν .
In the US, AM stations all have longer wavelengths than the FM stations.
•AM broadcast stations are licensed to operate only in the band 550-1700 kHz
•FM broadcast stations are licensed to operate only in the band 88-108 MHz
Visible Light
Microwaves
Radio/TV
Radar
Ultraviolet
Infrared
Gamma Rays
X-Rays
Low
High
Long
Short
Low
High
Energy
Red
Orange
Yellow
Green
Blue
Violet
Energy
• Incandescent light bulbs give off most of
their energy in the form of heat-carrying
infrared light photons -- only about 10
percent of the light produced is in the visible
spectrum. This wastes a lot of electricity.
Cool light sources, such as fluorescent
lamps and LEDs, don't waste a lot of energy
generating heat -- they give off mostly
visible light. For this reason, they are
slowly edging out the old reliable light bulb.
http://home.howstuffworks.com/light-bulb2.htm
Quantum Theory
A. Planck’s Hypothesis: (Max Planck
1900)
1.
2.
Studied emission of light from hot objects
Observed color of light varied with
temperature
3. Suggested the objects do not continuously
emit E, but emit E in small specific
amounts
a. Light is absorbed or emitted in a little
packet or bundle called a quantum
(quanta –plural).
b. Quantum = minimum amount of E
that can be lost or gained by an atom
c. Energies are quantized.
(Think steps not a ramp)
eeee-
X
Max Planck’s
Energy Equation
4. Proposed that energy is directly proportional
to frequency.
E = h
Planck’s equation for each quantum
h = Plank’s constant = 6.626 x 10-34 J.s
This is a direct relationship.
As energy increases, frequency increases.
Albert Einstein
While well-known for the equation
E=mc2 , Einstein’s work on the
photoelectric effect resulted in being
awarded the 1921 Nobel Prize in
Physics.
(1879 – 1955)
German Physicist
Albert Einstein and the
Photoelectric Effect
Refers to the emission of
electrons from a metal
when light shines on the
metal
-
Observations:
1. Electrons are ejected by light of
sufficient energy. Energy
minimum is different for
different metals.
2. The current (# of electrons
emitted/s) increases with
brightness of the light.
+
my.hrw.com
Albert Einstein and the
Photoelectric Effect
Conclusions:
1. Proposed that light consists of quanta of
energy that behaves like particles.
2. Quantum of light = photon = massless
particle that carries a quantum of energy.
3. Proposed the Dual Nature of Light: its
wave and particle nature.
a) Light travels through space as waves
b) Light acts as a stream of particles when it
interacts with matter.
my.hrw.com
Light (Electromagnetic Radiation)
Spectroscopy
Definition: a method of studying substances
that are exposed to some sort of continuous
exciting energy.
A. Emission Line Spectra: contains only certain colors
or wavelengths (  ) of light.
1. Every element has its own line spectrum (fingerprint).
Continuous Spectrum – White Light
Line Spectrum – Excited Elements
White light gives off a Continuous Spectrum
a blending of every possible wavelength
Gas Discharge Tubes
• Electricity is added to the gas which causes the electrons to
jump to a higher or excited state. They immediately fall
back to the ground state and give off particular wavelengths
of light. We see a blending of wavelengths without the
spectroscopes.
Flame Tests
• used to test
qualitatively for the
presence of certain
metals in chemical
compounds.
• the heat of the
Bunsen flame
excites electrons
that emit visible
light.
Copper(II) sulfate
Lithium chloride
Potassium chloride
Barium nitrate
Spectroscope
• Uses a diffraction grating to diffract the light
into particular wavelengths of light.
A Line Spectra result from excited elements - as
electrons of an element gain energy and rise to an
excited state they then fall back to their ground state in
the same pattern producing the same energy drop each
time which we see as individual wavelengths of light.
Atomic Spectra and the Bohr Model of
Hydrogen (1913)
Neils Bohr - Danish Scientist
Explained the bright-line spectrum of
hydrogen
Study:
• Added E as electricity to H gas at low
pressure in a tube.
• Emitted E as visible light, was
observed through a prism
Result: Hydrogen emitted 4
distinct bright lines of color,
aka bright line spectrum
Electrons release energy as they fall
back to a lower energy level
Electrons absorb energy to rise to a higher or excited
state and emit energy in the form of a photon of light
as they fall back to their ground states.
Path of an excited electron as it “falls”
back to the Ground State
• When electrons gain
energy, they jump to a
higher energy level
(excited state).
• Electrons are not stable at
the excited state and will
immediately fall back to a
lower level or ground state.
• As they fall, they emit
electromagnetic radiation.
• Depending on how far
they fall determines the
type of radiation (light)
released.
Bohr Model of Hydrogen
Conclusion:
• *Unique line spectrum is due to quantized electron energies.
• *Electrons are in specific orbits related to certain amounts of
energy known as stationary states.
• *Orbits are related to energy levels.
• *Energy levels are identified as E1, E2, E3, … (n = 1, 2, 3, …)
• *Lowest energy level = ground state
• *Electrons absorb certain amounts of energy to move to a higher
energy level farther away from the nucleus = excited state
• *Electrons return to the more stable ground state and release a
photon that has energy equal to the difference in energy between
the energy levels.
– from E2 to E1: Ephoton = E2 – E1 (difference in energy)
The Bohr Atom for Hydrogen a Model
1. Successful in calculating the wavelength,
frequency, & energy of hydrogen’s line
spectrum.
2. Successful in calculating the energy needed to
remove hydrogen’s electron
H(g) + energy  H+1(g) + 1eCalculated ionization E = observed ionization E
= 1312.1 kJ/mol
Lyman, Balmer and Paschen series
of the Hydrogen Atom
• Lyman series:
electrons fall to n = 1
and give off UV light.
• Balmer series:
electrons fall to n = 2
and give off visible
light.
• Paschen series:
electrons fall to n = 3
and give off infrared
light.
When electrons absorb energy they jump to a higher
(excited) state.
n=2 n=3 n=4 n=5 n=6 n=7
Electrons are not stable. Radiation (light) is emitted when
an electron falls back from a higher level to a lower level.
Infrared Light
Visible Light
Ultraviolet Light
Atomic Spectra
Hydrogen
Helium
Lithium
Mercury
Although Bohr’s
atomic model
explained the line
spectra of hydrogen,
it failed for heavier
elements.
Limitations of the Bohr Model
a. Model could not calculate the
wavelengths of observed
spectra of multi-electron
atoms.
b. Model could not explain the
chemical behavior of atoms.
c. Bohr used classical mechanics
to understand the behaviors of
small particles.
d. The Bohr model is also known
as the planetary, solar system,
or satellite model.
Quantum Mechanical Model
of the Atom
A. Louie De Broglie (1924-5)
1. Took Einstein’s idea that light can
exhibit both wave and particle
properties
2. Very small particles (like electrons)
display properties of waves.
3. Behavior of electrons in Bohr’s
quantized orbits was similar to behavior
of waves
French scientist
Known: any wave confined to a space can only have
specific frequencies
De Broglie suggested electrons are waves confined to the
space around the atomic nucleus.
Electrons could exist only at specific frequencies which
correspond to specific energies (E = h quantized E of
Bohr)
Quantum Mechanical Model
of the Atom
4. Experimentally proven in 1927 by diffraction of
electrons by Davisson & Germer (showed
diffraction of electrons by a crystal of Ni)
B. Wave-Particle Duality of Nature
a. Light and electrons (very small particles like
electrons, atoms, molecules) have properties of
waves and particles  QUANTUM
MECHANICS (based on WAVE properties)
**Large objects obey the laws of classical mechanics**
C. Werner Heisenberg: (1927)
1. Heisenberg’s Uncertainty Principle:
states that it is impossible to determine
simultaneously both the position and
velocity of an electron or any other
particle.
2. You cannot predict future locations of
particles.
3. He found a problem with the Bohr Atom
- no way to observe or measure the orbit
of an electron.
D. Erwin Schrödinger
Wave Equation (1926)
1. Wave nature of an electron is described by a
mathematical equation.
2. Four quantum numbers in the equation are
used to describe an electron’s behavior –
location and energy.
3. Electron is treated as a wave with
quantized energy.
4. Describes the probability of the
electrons found in certain
locations around the nucleus.
(1887 – 1961)
Austrian Physicist
Electron Density
An orbital is a region in
which an electron with a
particular energy is likely
to be found.
Where the density of an
electron cloud is high there
is a high probability that is
where the electron is
located. If the electron
density is low then there is
a low probability.
E. Atomic Orbitals - region around the
nucleus where an electron with a
particular energy is likely to be found
(not the same as Bohr’s orbits!)
1. Orbitals have characteristic shapes, sizes, &
energies.
2. Orbitals do not describe how the electron moves.
3. The drawing of an orbital represents the
3-dimentional surface within which the electron
is found 90% of the time.
4. Sublevels can have 4 different shapes
s – orbital spherical
1s, 2s & 3s orbitals Superimposed on
one another
Electron-Cloud Models
p-orbital – dumbbell shaped
p-orbital - dumbbell shaped
d-orbital - double dumbbell or fan blades
s,p and d orbitals
z
z
z
x
x
x
y
y
x
y
s orbital
y
p orbitals
z
z
z
x
x
y
z
y
z
x
y
z
x
y
x
y
d orbitals
For a more complete representation and presentation of atomic orbitals go to
http://winter.group.shef.ac.uk/orbitron/
Models of d-orbitals
f-orbital –
more complex!
f orbitals
f – orbitals (3D)
Quantum Numbers
• Each quantum number provides more specific
information on the probable location of an
electron.
• Each electron within an atom can be described
by a unique set of 4 quantum numbers.
Quantum Numbers - Finding an
address for each electron:
1. “state” Principle Quantum Number (n) or the
energy level;
a. Describes the relative size of the electron cloud.
b. Positive integer values (n = 1 to n = 7)
2. “city”
a.
b.
c.
d.
Sublevel (l)
Describes the shape of the electron cloud.
The maximum number of sublevels within a level = n
Shapes are s, p, d,or f.
Lowest energy = s
Highest energy = f
Quantum numbers cont.
3. “street” Orbital (ml) odd # of orbitals
1. Describes the orientation or direction in space
a) s – 1 orbital
b) p – 3 orbitals (x, y, z)
c) d – 5 orbitals (xy, yz, xz, x2 – y2, z2)
d) f – 7 orbitals (y3 – 3yx2, 5yz2-yr2, x3-3xy2,
zx2-zy2, xyz, 5xz2-3xr2, 5z3-3zr2)
2. Orbitals within the same sublevel have the same
energy are called degenerate orbitals
3. An orbital can hold a maximum of 2 electrons
Quantum numbers cont.
4. “house” Spin (ms)
1. Describes the direction of electron spin in an
orbital.
2. The clockwise or counterclockwise motion of
electrons.
3. Only electrons with opposite spins can occupy the
same orbital.
4. The opposite spin is written as
+1/2 or -1/2 or  or
E. Electron Configurations:
1. Shorthand notation for indicating the number of
electrons in each level, sublevel, and orbital.
1s2
2. Shows the distribution of electrons among the
orbitals. Describes where the electrons are found
& what energy they possess.
Electron Configuration Rules
1. The Aufbau
Principle:
electrons are
added one at a
time to the
lowest energy
orbital
available.
Pauli Exclusion Principle:
1. Each orbital can only hold 2 electrons.
2. The electrons must have opposite spins.
s-sublevel
p-sublevel
d-sublevel
f-sublevel
=
=
=
=
max 2 electrons
max 6 electrons
max 10 electrons
max 14 electrons
incorrect: ↑↑↑ incorrect: ↑↑ correct: ↑↓
Hund’s Rule:
• Electrons will remain
unpaired in
degenerate orbitals
before they pair up.
incorrect ↑↓ ↑ __
correct
↑ ↑ ↑
Electron Blocks on the
Periodic Table
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
3p
3s
2p
2s
1s
6d
5d
4d
5f
4f
3d
Pauli Exclusion Principle: No more
than 2 e- are put in each orbital and
they must have opposite spin.
Hund’s Rule: electrons spread out
among equal energy orbitals in a
sublevel (like charges repel)
Aufbau Principle: Electrons fill
lowest energy levels first (n=1)
Electron Configuration Examples:
Ex) electron configuration for Na:
1s2 2s2 2p6 3s1
Ex) orbital filling box diagram for Na:
x
y
z
      _
1s
2s
2p
3s
Electron Dot Diagrams:
Write the symbol for the element.
Place dots around the symbol to represent the
valence s & p electrons only.
Do NOT include d & f orbitals in diagram.
p orbital electrons
s orbital electrons
Electron Configuration
Orbital Box Diagram
1s22s22p4
O
z
    
1s
2s
2p
x
35
17
y
x
16
8
Electron-dot Diagram
22s22p63s23p5
1s
Cl
y
z
x
y
z
        
1s
2s
2p
3s
3p
1s22s22p63s23p64s23d104p65s24d105p4
127
52
Te
x
y
z
x
y
z
x
y
z
x
y
z
                          
1s
2s
2p
3s
3p
4s
3d
4p
5s
4d
What does the Tellurium electron-dot resemble???
5p
Mark your Periodic Tables
1
2
13
14
15
16
17
18
Unpaired vs. Paired Electrons
Filled and Half-filled orbitals
• Atoms with unpaired electrons are said to be
paramagnetic. These are weakly attracted to a
magnetic field.
• Atoms with all paired electrons are said to be
diamagnetic. These are weakly repelled from a
magnetic field.
• ½ filled and filled orbitals have special stability
Noble Gas or Shorthand
Electron Configurations
• Rb
1
[Kr]5s
• Se
2
10
[Ar]4s 3d 4p
4
• At
2
14
10
[Xe]6s 4f 5d 6p
5
Draw the Dot Diagrams for these elements
Exceptions to the Rules
• Max stability - ½ filled and
filled orbitals
–Cr
–Mo
–Cu
–Ag
–Au
[Ar]4s2 3d 4
2
[Ar]4s 3d
9
[Ar]4s1 3d5
1
10
[Ar]4s 3d
Exceptions to the Rules
• Max stability - ½ filled and
filled orbitals
–Cr
–Mo
–Cu
–Ag
–Au
[Ar]4s2 3d 4
1
5
2
9
1
10
[Kr]5s 4d
[Ar]4s 3d
[Kr]5s 4d
[Xe]6s1 4f 14 5d10
[Ar]4s1 3d5
1
10
[Ar]4s 3d
Electron Configuration
for Ions
• K
[Ar]
[Ar]4s1
1s 2 2s 2 2p6 3s 2 3p6 4s1
• K+1 1s2 2s2 2p6 3s2 3p6
• P
[Ne]3s 2 3p3
1s 2 2s 2 2p6 3s 2 3p3
• P-3
• Al
[Ne]3s2 3p1
1s2 2s2 2p6 3s2 3p1
• Al+3
• Se
2
10
[Ar]4s 3d 4p
4
1s2 2s2 2p6 3s 2 3p6 4s 2 3d10 4p 4
[Ne]3s 2 3p6
1s2 2s2 2p6 3s 2 3p6
[Ne]
2
2
1s 2s 2p
6
• Se-2 [Ar]4s2 3d10 4p6
1s2 2s2 2p6 3s 2 3p6 4s 2 3d10 4p6
Excited vs. Ground State
• If an electron absorbs energy,
it is in an EXCITED state
Ne: 1s22s22p53s1
• How is this different from the
ground state configuration?
Ne: 1s22s22p6