Inorganic Pharmaceutical Chemistry
Lecture No. 3
Date :18/10 /2012
Dr. Mohammed Hamed
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Properties of Monatomic Ions
The electrons in the outermost shell (the ones with the highest value of n) are
the most energetic, and are the ones which are exposed to other atoms. This
shell is known as the valence shell. The inner, core electrons (inner shell) do not
usually play a role in chemical bonding.
Elements with similar properties generally have similar outer shell
configurations. For instance, we already know that the alkali metals (Group I)
always form ions with a +1 charge; the "extra" s1 electron is the one that's lost:
IA Li 1s22s1
Li+
1s2
Na 1s22s22p63s1
Na+ 1s22s22p6
K 1s22s22p63s23p64s1
K+
1s22s22p63s23p6
The next shell down is now the outermost shell, which is now full — meaning
there is very little tendency to gain or lose more electrons. The ion's electron
configuration is the same as the nearest noble gas — the ion is said to be
isoelectronic with the nearest noble gas. Atoms "prefer" to have a filled
outermost shell because this is more electronically stable.
The Group IIA and IIIA metals also tend to lose all of their valence electrons to
form cations.
IIA Be 1s22s2
Mg 1s22s22p63s2
IIIA Al 1s22s22p63s23p1
Be2+ 1s2
Mg2+ 1s22s22p6
Al3+ 1s22s22p6

The Group IV and V metals can lose either the electrons from the p
subshell, or from both the s and p subshells, thus attaining a pseudo-noble
gas configuration.
IVA Sn [Kr]4d105s25p2
Sn2+ [Kr]4d105s2
Sn4+ [Kr]4d10
Pb [Xe]4f145d106s26p2
Pb2+ [Xe]4f145d106s2
Pb4+ [Xe]4f145d10
VA Bi [Xe]4f145d106s26p3


Bi3+
[Xe]4f145d106s2
Bi5+
[Xe]4f145d10
The Group IV - VII non-metals gain electrons until their valence shells are
full (8 electrons).
IVA C 1s22s22p2
C4-
1s22s22p6
VA N 1s22s22p3
N3-
1s22s22p6
VIA O 1s22s22p4
O2-
1s22s22p6
VIIA F 1s22s22p5
F-
1s22s22p6
The Group VIII noble gases already possess a full outer shell, so they have
no tendency to form ions.
VIIIA Ne 1s22s22p6
Ar 1s22s22p63s23p6

Transition metals (B-group) usually form +2 charges from losing the
valence s electrons, but can also lose electrons from the highest d level to
form other charges.
B-group Fe 1s22s22p63s23p63d64s2
Fe2+ 1s22s22p63s23p63d6
Fe3+ 1s22s22p63s23p63d5
Periodic table
A periodic table is a tabular display of the chemical elements, organized on the
basis of their atomic numbers, electron configurations, and recurring chemical
properties. Elements are presented in increasing atomic number. The main body
of the standard form of table is an 18 × 7 grid, and elements with the same
number of valence electrons are kept together in groups, such as the halogens
and the noble gases. There are four distinct rectangular areas or blocks. The fblock is usually not included in the main table, but rather is floated below, as an
inline f-block would often make the table impractically wide. Using periodic
trends, the periodic table can help predict the properties of various elements and
the relations between properties. As a result it provides a useful framework for
analyzing chemical behavior, and is widely used in chemistry and other
sciences.
Although precursors exist, Dmitri Mendeleev is generally credited with the
publication, in 1869, of the first widely recognized periodic table. He developed
his table to illustrate periodic trends in the properties of the then-known
elements. Mendeleev also predicted some properties of then-unknown elements
that would be expected to fill gaps in this table. Most of his predictions were
proved correct when the elements in question were subsequently discovered.
Mendeleev's periodic table has since been expanded and refined with the
discovery or synthesis of further new elements and the development of new
theoretical models to explain chemical behavior.
All elements from atomic numbers 1 (hydrogen) to 118 (ununoctium) have been
discovered or synthesized. Of these, all up to and including californium exist
naturally; the rest have only been artificially synthesised in laboratories.
Production of elements beyond ununoctium is being pursued, with the question
of how the periodic table may need to be modified to an extended form to
accommodate these elements being a matter of ongoing debate. Numerous
synthetic radionuclides of naturally occurring elements have also been produced
in laboratories.
In most modern periodic tables, the elements are placed progressively in each
period from left to right in the sequence of their atomic numbers, with a new
row started after a noble gas. The first element in the next row is always an
alkali metal with an atomic number one greater than that of the noble gas (e.g.
after krypton, a noble gas with the atomic number 36, a new row is started by
rubidium, an alkali metal with the atomic number 37). No gaps currently exist
because all elements between hydrogen and ununoctium (element 118) have
been discovered. Since the elements are sequenced by atomic number, sets of
elements are sometimes specified by terms such as "through" (e.g. through iron),
"beyond" (e.g. beyond uranium), or "from ... through" (e.g. from lanthanum
through lutetium). The terms "light" and "heavy" are sometimes also used
informally to indicate relative atomic numbers, as in "lighter than carbon" or
"heavier than lead", although technically the weight or mass of atoms of an
element (their atomic weights or atomic masses) do not always increase
monotonically with their atomic numbers. For instance tellurium, element 52, is
on average more massive than iodine, element 53. More often, however,
elements are referred to as light or heavy on account of their densities.
The significance of atomic numbers to the organization of the periodic table was
not appreciated until the existence and properties of protons and neutrons
became understood. Mendeleev's periodic tables used atomic weight instead of
atomic number to organize the elements, information determinable to fair
precision in his time, which worked well enough in most cases to give a
presentation that was able to predict other elements' properties far better than
any other method known at that time. Substitution of atomic numbers, once
understood, gave a definitive, integer-based sequence for the elements, still used
today even as new synthetic elements are being produced and studied.
Groups
A group or family is a vertical column in the periodic table. Groups usually have
more significant periodic trends than periods and blocks, In some groups, the
elements have very similar properties and exhibit a clear trend in properties
down the group.
Under an international naming convention, the groups are numbered
numerically from 1 to 18 from the leftmost column (the alkali metals) to the
rightmost column (the noble gases). Previously, the groups were known by
roman numerals. In America, the roman numerals were followed by either an
"A" if the group was in the s- or p-block, or a "B" if the group was in the dblock. The roman numerals used correspond to the last digit of today's naming
convention (e.g. the group 4 elements were group IVB, and the carbon group
were group IVA). In Europe, the lettering was similar, except that "A" was used
if the group was before group 10, and "B" was used for groups including and
after group 10. In addition, groups 8, 9 and 10 used to be treated as one triplesized group, known collectively in both notations as group VIII. In 1988, the
new IUPAC naming system was put into use, and the old group names were
deprecated.
Elements in the same group show patterns in atomic radius, ionization energy,
and electronegativity. From top to bottom in a group, the atomic radii of the
elements increase. Since there are more filled energy levels, valence electrons
are found farther from the nucleus. From the top, each successive element has a
lower ionization energy because it is easier to remove an electron since the
atoms are less tightly bound. Similarly, a group has a top to bottom decrease in
electronegativity due to an increasing distance between valence electrons and
the nucleus. There are exceptions to these trends, however, an example of which
is that in group 11, the electronegativity increases farther down the group.
Periods
A period is a horizontal row in the periodic table. Although groups generally
have more significant periodic trends, there are regions where horizontal trends
are more significant than vertical group trends, such as the f-block, where the
lanthanides and actinides form two substantial horizontal series of elements.
Elements in the same period show trends in atomic radius, ionization energy,
electron affinity, and electronegativity. Moving left to right across a period,
atomic radius usually decreases. This occurs because each successive element
has an added proton and electron which causes the electron to be drawn closer to
the nucleus. This decrease in atomic radius also causes the ionization energy to
increase when moving from left to right across a period. The more tightly bound
an element is, the more energy is required to remove an electron.
Electronegativity increases in the same manner as ionization energy because of
the pull exerted on the electrons by the nucleus. Electron affinity also shows a
slight trend across a period. Metals (left side of a period) generally have a lower
electron affinity than nonmetals (right side of a period) with the exception of the
noble gases.
Blocks
Because of the importance of the outermost electron shell, the different regions
of the periodic table are sometimes referred to as blocks, named according to the
subshell in which the "last" electron resides. The s-block comprises the first two
groups (alkali metals and alkaline earth metals) as well as hydrogen and helium.
The p-block comprises the last six groups which are groups 13 to 18 in IUPAC
(3A to 8A in American) and contains, among others, all of the metalloids. The
d-block comprises groups 3 to 12 in IUPAC (or 3B to 2B in American group
numbering) and contains all of the transition metals. The f-block, usually offset
below the rest of the periodic table, comprises the lanthanides and actinides.
Variations and other conventions
In presentations of the periodic table, the lanthanides and the actinides are
customarily shown as two additional rows below the main body of the table,
with placeholders or else a selected single element of each series (either
lanthanum or lutetium, and either actinium or lawrencium, respectively) shown
in a single cell of the main table, between barium and hafnium, and radium and
rutherfordium, respectively. This convention is entirely a matter of aesthetics
and formatting practicality; a rarely used wide-formatted periodic table inserts
the lanthanide and actinide series in their proper places, as parts of the table's
sixth and seventh rows (periods).
Many presentations of the periodic table show a dark stair-step diagonal line
along the metalloids, with metals to the left of the line and non-metals to the
right. Various other groupings of the chemical elements are sometimes also
highlighted on a periodic table, such as transition metals, post-transition metals,
and metalloids. Other informal groupings of the elements exist, such as the
platinum group and the noble metals, but are rarely addressed in periodic tables.
Effective Nuclear Charge
The effective nuclear charge is the net positive charge experienced by
an electron in a multi-electron atom. The term "effective" is used because
the shielding effect of negatively charged electrons prevents higher orbital
electrons from experiencing the full nuclear charge by the repelling effect of
inner-layer electrons. The effective nuclear charge experienced by the outer
shell electron is also called the core charge. It is possible to determine the
strength of the nuclear charge by looking at the oxidation number of the atom
Calculating the effective nuclear charge
In an atom with one electron, that electron experiences the full charge of the
positive nucleus. In this case, the effective nuclear charge can be calculated
from Coulomb's law.
However, in an atom with many electrons the outer electrons are simultaneously
attracted to the positive nucleus and repelled by the negatively charged
electrons. The effective nuclear charge on such an electron is given by the
following equation:
Zeff = Z − S
where
Z is the number of protons in the nucleus (atomic number), and
S is the average number of electrons between the nucleus and the electron
in question (the number of nonvalence electrons).
S can be found by the systematic application of various rule sets, the
simplest of which is known as "Slater's rules".
Note: Zeff is also often written Z*.
Values Shielding effect
The shielding effect describes the decrease in attraction between an electron and
the nucleus in any atom with more than one electron shell. It is also referred to
as the screening effect or atomic shielding.
Slater's rules
In quantum chemistry, Slater's rules provide numerical values for the effective
nuclear charge concept. In a many-electron atom, each electron is said to
experience less than the actual charge owning to shielding or screening by
the other electrons. For each electron in an atom, Slater's rules provide a value
for the screening constant, denoted by s, S, or σ, which relates the effective and
actual nuclear charges as
Rules
Firstly, the electrons are arranged in to a sequence of groups in order of
increasing principal quantum number n, and for equal n in order of
increasing azimuthal quantum number l, except that s- and p- orbitals are
kept together.
[1s] [2s,2p] [3s,3p] [3d] [4s,4p] [4d] [4f] [5s, 5p] [5d] etc.
Each group is given a different shielding constant which depends upon the
number and types of electrons in those groups preceding it.
The shielding constant for each group is formed as the sum of the
following contributions:
1. An amount of 0.35 from each other electron within the same group
except for the [1s] group, where the other electron contributes only
0.30.
2. If the group is of the [s p] type, an amount of 0.85 from each
electron with principal quantum number (n) one less and an amount
of 1.00 for each electron with an even smaller principal quantum
number
3. If the group is of the [d] or [f], type, an amount of 1.00 for each
electron inside it. This includes i) electrons with a smaller principal
quantum number and ii) electrons with an equal principal quantum
number and a smaller azimuthal quantum number (l)