Download The Development of the Periodic Table

The Periodic Table
The Development of the Periodic Table
• In the time of Mendeleev, we had not yet discovered the electron
• Mendeleev was convinced that the properties of elements had a periodicity if
we classified the elements in order of increasing atomic weight
• He had so much confidence in the periodicity of properties that he
suggested that some elements remained to be discovered, and he left
"holes" in his table to accommodate them once they were discovered
• He also predicted the properties of these remaining elements very
well
• One problem was the fact that Ar (element 18) had a greater mass than Na
(element 19)
• After the discovery of the nucleus by Rutherford, it was found that
the number of protons in the nuclei was the important criterion (not
the mass)
The Periodic Classification of Elements
• Each column (group) of the modern periodic table shares the same electronic
configuration for their valence electrons
• There are a few exceptions for the transition metals, the lanthanides, and the
actinides
The Periodic Classification of Elements
• The valence electrons are the peripheral electrons (the electrons with the
highest principal quantum number) of an atom
• Valence electrons are those involved in the formation of chemical bonds
• The fact that each member of a group shares the same number of valence
electrons explains the similarities in their reactivity
• N.B. The properties of the elements in IA, IIA, and VIIA groups are very
similar through the group
• In the IIIA, IVA, VA, et VIA groups, the properties of the elements
sometimes change gradually when descending the group
The Electronic Configurations of Cations and Anions
• In an ionic compound, the cation of a representative element is produced by
releasing electrons so that the cation reaches the electron configuration of a
noble gas:
eg.; Na : [Ne]3s1
devient Na  : [Ne]
becomes
Ca : [Ne]3s 2
devient Ca 2 : [Ne]
becomes
Al : [Ne]3s 2 3p1 becomes
devient Al 3 : [Ne]
• In an ionic compound, the anion of a representative element is produced by
accepting electrons so that the anion reaches the electron configuration of a
noble gas:
becomes
eg.; H : 1s1
devient H  : 1s 2 ou [He]
F : [He]2s 2 2p 5 becomes
devient F  : [He]2s 2 2p 6 ou
or [Ne]
[Ne]
or [Ne]
[Ne]
O : [He]2s 2 2p 4 becomes
devient O 2 : [He]2s 2 2p 6 ou
N : [He]2s 2 2p 3
devient N 3 : [He]2s 2 2p 6 ou
or[Ne]
[Ne]
becomes
The Electronic Configurations of Cations and Anions
• These ions of representative elements and noble gases are isoelectronic (have
the same number of electrons and therefore share the same electronic
configuration)
• For transition metals, we can often find more than one type of cations and
these cations are often not isoelectronic with a noble gas
• ex.; neither Fe2+ nor Fe3+ are isoelectronic with a noble gas, but each
can be found in nature
• Even if the ns orbital is filled before the (n-1)d orbital, when a transition metal
forms a cation, it is the ns orbital that empties first
• The ns and (n-1)d orbitals are very close in energy, and the electronelectron and electron-nucleus change when going from atom to cation,
thus causing a change in the energetic order between ns and (n-1)d
• ex.; Mn2+ has the configuration [Ar]3d5 and Zn2+ has the configuration
[Ar]3d10
The Effective Nuclear Charge
• The effective nuclear charge, Zeff, is the positive charge felt by a valence
electron
• Zeff  Z (where Z is the nuclear charge) since the other electrons screen the
nucleus from the valence electrons
• Zeff = Z -  where  is the screening constant
• To give an example of the importance of this screening effect, consider the
two ionization energies of He
• The first electron is extracted with 2373 kJ of energy
• It takes 5251 kJ to remove the second electron
• Without the screening effect, the two ionization energies would be the
same
The Effective Nuclear Charge
• All of the electrons contribute to the screening effects, but in the simplest
model,  is the number of electrons in the layers with principal quantum
numbers lower than that of the valence electrons
• ex.; Mg: 1s22s22p63s2
• ex.; Mg2+: 1s22s22p6
Z = 12,  = 10, Zeff = +2
Z = 12,  = 2, Zeff = +10
• N.B. These effective nuclear charges are what are felt by the valence
electrons (n = 3 for the atom, et n = 2 for the cation, in these
examples)
• For a neutral atom of an element, the effective nuclear charge is the number
of the element’s group (according to this simple model)
The Atomic Radius
• The radius is a poorly defined property
because the electron density, in principle,
extends to infinity
• The size of an atom can be estimated from
experimental data
– in a metal, the atomic radius is more or
less half the distance between two
adjacent atoms
– for an element that exists as a diatomic
molecule, the atomic radius is more or
less half the length of the bond
The Atomic Radius
• Going from left to right, the
atomic radii of elements decrease
• This decrease in the atomic radii
is due to the increasing effective
nuclear charge as we go from left
to right
• valence electrons are held
closer to the nucleus by
the increasing effective
nuclear charge
The Atomic Radius
• Descending a group, the atomic
radius increases
• When descending a group, the
effective nuclear charge seen by
the valence electrons remains more
or less the same
• However, the valence electrons
occupy a higher principal quantum
level, and thus the average
distance from the nucleus
increases
The Ionic Radius
• The ionic radius is the radius of a cation
or an anion
• When an atom becomes an anion, its size
increases because of the greater repulsion
between the electrons, thus causing the
electrons to occupy a larger volume
• When an atom becomes a cation, the
opposite occurs, and the radius decreases
• In addition, in many cases, the
valence shell with quantum level
n is emptied and the valence shell
then becomes the one with
principal quantum number (n-1)
The Atomic Radius
• Even if N3-, O2-, F-, Na+, Ca2+, et Al3+
are isoelectronic, their ionic radii are
very different due to their effective
nuclear charges
• Their configuration is 1s22s22p6 and
=2, and therefore Zeff = +5, +6, +7,
+9, +10, and +11, respectively, going
from N3- to Al3+
• A greater effective nuclear charge
more strongly attracts the electrons
towards the nucleus and the ionic
radius gets smaller
The Atomic Radius
• N.B. Atomic and ionic radii may differ substantially
Atomic and Ionic Radii
• Example: Arrange the following atoms in order of decreasing radii :
C, Li, Be.
• Solution: In general, the radius decreases from left to right, thus Li >
Be > C
• Example: Decide which ion is the smallest in each of the following
pairs : (a) K+, Li+ ; (b) Au+, Au3+ ; (c) P3-, N3• Solution:
(a) Li+, because the filled quantum level is n=1 rather than n=3 for K+
(b) Au3+ because there will be less repulsion between electrons
(c) N3- because the filled quantum level is n=2 rather than n=3 for P3-
Ionization Energy
• The ionization energy is the minimum energy required to extract an electron
from an isolated atom in its ground state
energy + X(g)  X+(g) + e• The higher the ionization energy, the more difficult it is to remove an electron
from the atom
• There are also second and third, …., ionization energies
energy + X+(g)  X2+(g) + eenergy + X2+(g)  X3+(g) + e….
• It becomes more difficult to extract the second, third , ... . , electron since the
repulsion between electrons decreases after each ionization and it is very
difficult to separate an electron from a cation due to the favorable electrostatic
interaction between these two opposite charges
Ionization Energy
• The ionization energy decreases when
descending a group
• The effective nuclear charge does not
change, but the valence electron that is
removed is further from the core, and
therefore more weakly retained and
easily removed
• The ionization energy increases, in general,
when going from left to right on the periodic
table
• From left to right, the effective nuclear
charge increases and the valence
electron that is removed is more
strongly held by the nucleus
Ionization Energy
• There are two exceptions when going from
left to right
• A group IIIA (ns2np1) element is easier to
ionize than a group IIA (ns2) element
• For a group IIIA element, an
electron is removed from the np
orbital which is higher in energy
than the ns orbital of a group IIA
element
• The increase in effective charge
once again becomes the dominant
factor when we reach group IVA
Ionization Energy
• A group VIA (ns2np4) element is easier to
ionize than a group VA (ns2np3) element
• For a group VIA element, there are
two electrons in one of the np
orbitals, and these two electrons
repel each other, and one of the two
is more easily removed than the np
electrons of an element in Group VA
where each np electron has an orbital
to itself
• The increase in effective charge
once again becomes the dominant
factor when we reach group VIIA
The Unique Characteristics of the Second Period
• The elements of the second period are Li to F
• The properties of the first member of each representative group (the
elements of the second period) are often very different from other
group members, i.e., the first member is “unique”
• eg.; The chemistry of Si resembles the chemistry of Ge more
then it does the chemistry of C
• eg.; The ability to form good double bonds is limited almost
exclusively to the atoms of the second period
• This effect is mainly due to the relatively small atomic radius
of the first group member
The Diagonal Relationship
• The diagonal relationship refers to
similarities between two immediately
adjacent elements of different groups
and different periods on the periodic
table (ie., diagonally)
• The effect is most important for Li/Mg,
Be/Al, and B/Si
• The effect is largely due to the fact that
the charge/radius ratio of their cations
are very similar