PERIODIC TABLE
History of the periodic table
Johann Dobereiner (1780-1849).
•first known chemist to recognize periodic
patterns.
•suggested a “triad” system for
classification, but found it limiting.
•John Newlands.
•* In 1864, he suggested that elements
should be arranged in “octaves,” based on
the idea that certain properties seemed to
repeat every eighth element; still not very
successful.
•.
Atomic Mass
• Avogadro had suggested the idea of a mole.
• Which was a specific number of atoms, he had no idea
of actual the amount, 6.022x1023.
• He said equal volumes of gases at the same pressure
and temperature would have the same number of
particles, thus 1 mole of any gas at STP = 22.4 L
• By reacting moles of elements with each other you can
determine what one mole of every element is.
• If you take the mass of one mole of an element you
have the atomic mass.
Mendeleev and Meyer
•Julius Meyer and Dmitri Mendeleev both
independently conceived the present
periodic table.
•Mendeleev is given credit because of his
ability to predict accurately the
properties of undiscovered elements.
Dmitri Mendeleev
• Element #101 was named in his honor
• first to come up with a periodic table
• ~1870 there were 63 elements known to
man, he organized the elements in order of
their atomic mass, he saw a pattern form
with the # of elements that can be bonded
to that element.
More Mendeleev
• He arranged his table so that columns of elements with
repeating patterns formed, he left spaces in his table
where it appeared an element wasn’t discovered yet
• He then predicted the properties of these elements
by looking at the other elements in the family and period
of that element
Medeleev’s Table
It may not sound like much, but…
• He predicted what undiscovered elements
would be like giving specifics!!!
• And he was Right!
• This is the basis for acceptance of a scientific
theory.
• If the model is correct then the scientist can
assume what would happen if…
1871 Discovery
Property
atomic mass
Ekaaluminum
(Mendeleev’s
prediction)
68
Gallium
69.72
density (g/cm³) 6.0
5.904
melting point
(°C)
29.78
Low
oxide's formula Ea2O3 (density - 5.5 g
-3
cm )
chloride's
formula
Ga2O3 (density - 5.88 g
-3
cm )
(soluble in both alkalis and
acids)
(soluble in both alkalis and acids)
Ea2Cl6 (volatile)
Ga2Cl6 (volatile)
1882 Discovery
Property
atomic mass
density (g/cm³)
melting point (°C)
color
Ekasilicon
72
5.5
high
grey
Germanium
72.61
5.35
947
grey
oxide type
oxide density (g/cm³)
oxide activity
refractory dioxide
4.7
feebly basic
refractory dioxide
4.7
feebly basic
chloride boiling point
chloride density (g/cm³)
under 100°C
1.9
86°C (GeCl4)
1.9
Dmitri Mendeleev
Problem with Mendeleev’s table
• Mendeleev arranged his table by atomic
mass
• a few elements appeared to be slightly out
of place, Mendeleev put them in the right
place and guessed that their atomic
masses were incorrectly measured
• Actually, he was arranging them by the
wrong number.
Atomic Number and Henry Moseley
• Moseley discovered that the charge of a
nucleus is always divisible by the same
number (which we now call of the charge of a
proton).
• This gave him an integral number for each
atom he called the atomic number.
• He corrected the periodic law when he stated
that periodicity is a result of increasing atomic
number, not increasing atomic mass.
Henry Moseley
Aufbau Principle
• As protons are added one by one to the
nucleus to build up the elements, electrons
are similarly added to hydrogen like orbitals
• The lowest energy level is the 1 s orbital
• From this we can begin to diagram the
orbitals in an atom
nucleus
•1 s __ electrons
+
•2 s __ p __ __ __
•3 s __ p __ __ __ d __ __ __ __ __
Aufbau exceptions
• You see the first exceptions to the Aufbau
principle in the 3d orbital.
• Chromium is [Ar]3d54s1
• As opposed to the expected [Ar]3d44s2
• It then goes back to “normal”, until copper.
• Copper is [Ar]3d104s1
• There are more examples is the d orbitals
below, and in the f orbitals
Hund’s Rule
• The lowest energy configuration (ground
state) for an atom is the one having the
maximum number of unpaired electrons
allowed by the Pauli Exclusion Principle
(two electrons per orbital).
• The convention is to represent the spins of
electrons as up or down arrows
Periodic Table
• The modern design of the periodic table is
built off of these orbital diagrams
You can determine the orbital by the
placement
H
For these elements the last
electrons are filling s orbitals
He
Li
For these elements the last
electrons are filling d orbitals
for these elements
the last electrons are filling
p orbitals
For the elements the last electrons
are filling f orbitals
Electrons configurations are written to
shorten orbital diagrams
• The electron configuration for potassium is
1s22s2p63s2p64s1
• This can be abbreviated to [Ar]4s1
• This is the same as saying everything that
is in Ar, 1s22s2p63s2p6 +4s1
Total order
1
• 1 s __
3
2
• 2 s __
p __ __ __
4
5
• 3 s __
p __ __
__ d
6
• 4 s __
7
__ __ __ __ __
8
p __ __
__ d __10__ __ __ __ f __ __13__ __ __ __ __
9
• 5 s __
p __ 11
__ __ d __14__ __ __ __ f __ __17__ __ __ __ __
12 p __ 15
• 6 s __
__ __ d __18__ __ __ __
16 p __ 19
• 7 s __
__ __
Aufbau Principle is the order of filling the
orbitals
• 4s is filled before 3d
• The (n+1)s orbitals always fill before the n d
orbitals, explained by the penetration
effect.
• The penetration effect is looking at the
forces on the interior electrons in an atom.
Not only is there an attraction to the
nucleus, but the outer electrons repulsion is
pushing these orbitals closer to the
nucleus, making them smaller.
Lone electron in the 5d or 6 d
• Another exception
• Lanthanum’s last electron fills the 5d orbital.
[Xe]6s25d1
• Ce has its last electron fill in the 4f, and the
one from 5d (La) jumps to 4f. [Xe]6s24f2
• The next elements all fill up the 4f until
Gadolinium which fills up the 5d again.
• A similar “jumping” of one electron occurs
with the actinide series
Different Periodic Tables
• The electron “jumping” is why there are two
forms of the periodic table in common
usage.
• One has lanthanum and actinium in with
the d filling elements.
• The other has lanthanum and actinium with
the f filling elements
To Determine the Condensed
electron configuration
• Start at the previous noble gas (lower
atomic number)
• Write out every element in order of
atomic number to your given
• Determine which shell all elements are
filling. The number of elements in that
shell is your superscript number.
For Iron
• Previous Noble Gas----Ar
• Write all elements
• K Ca Sc Ti V Cr Mn Fe
• These fill 4 s
• These fill 3 d
• So it is
• [Ar]3d64s2
Practice
• Write out the condensed electron
configuration and orbital diagram for…
• Nitrogen
• Tin
• Molybdenum
• Rutherfordium
Effective Nuclear Charge (Zeff)
•Zeff is the net positive charge experienced
by a specific electron in a multi electron
atom.
• As electrons get further away, they will be
“shielded” from the nucleus by other
electrons.
• What is the period trend, with explanation?
Ionization Energy
• The energy required to remove an electron.
•What type of ion is produced?
•Is the process endothermic or exothermic?
•What is the group trend, with explanation?
•What is the period trend, with explanation?
Exceptions
• Normally moving left to right across a period
will increase in ionization energy because it is
being held in by a stronger nucleus.
• However there is an observed dip in ionization
energy going from Be to B.
• Why?
• It is harder to remove an electron from 2p
than 3s, these aren’t hydrogen orbitals.
• This trend continues for Mg-Al, then you see
the trend for Zn-Ga, and Cd-In, and Hg-Tl
Another exception
• There is another between group 15 and 16
(nitrogen group to oxygen group).
• This exception is because in group 15 you
have one electron in each of the 3 p
orbitals, group 16 has a 4th electron that is
easier to remove.
Successive Ionization Energies
•Ionization energy doesn’t have to stop at one
electron, you can remove multiple electrons, to
get a 2nd ionization energy I2, 3rd ionization
energy I3 …
•Why are second ionization energies more than
first ionization energies?
•Electrons are more easily removed until what
occurs?
Trends in Ionization Energy
•The first ionization energy for
phosphorus is 1060 kJ/mol, and that for
sulfur is 1005 kJ/mol. Why?
•Consider atoms with the following
electron configurations:
•
1s22s22p6
•
1s22s22p63s1
•
1s22s22p63s2
• Which atom has the largest first
ionization energy, and which one has the
smallest second ionization energy?
Explain your choices.
Atomic Radius
• ~relative size of an atom.
• As you move down a period you are
increasing the number of energy levels so
the atom is getting larger.
• As you move left to right across a
period, a more positive nucleus pulls
everything in closer, making the atom
smaller.
• It increases down and to the left.
Electron Affinity
•Electron affinity is the change in energy
when an atom captures a bonding electron.
•Energy is normally released, so the number
is normally reported as a negative.
•What is the group trend, with explanation?
•What is the period trend, with
explanation?