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.
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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*.
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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 :
Zeff = Z - S
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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 0.35.
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) .
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Periodic properties :
Atomic Radius :
The atomic radius of an element is half of the distance between the
centers of two atoms of that element that are just touching each other.
Generally, the atomic radius decreases across a period from left to right
and increases down a given group. The atoms with the largest atomic
radii are located in Group I and at the bottom of groups. Moving from
left to right across a period, electrons are added one at a time to the
outer energy shell. Electrons within a shell cannot shield each other
from the attraction to protons. Since the number of protons is also
increasing, the effective nuclear charge increases across a period. This
causes the atomic radius to decrease.
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Moving down a group in the periodic table, the number of electrons and filled
electron shells increases, but the number of valence electrons remains the same.
The outermost electrons in a group are exposed to the same effective nuclear
charge, but electrons are found farther from the nucleus as the number of filled
energy shells increases. Therefore, the atomic radii increase.
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Ionization energy :
The ionization energy, or ionization potential, is the energy required to completely
remove an electron from a gaseous atom or ion .
The first ionization energy is the energy required to remove one electron from the
parent atom.
The second ionization energy is the energy required to remove a second valence
electron from the univalent ion to form the divalent ion.
The second ionization energy is always greater than the first ionization energy.
Ionization energies increase moving from left to right across a period (decreasing
atomic radius).
Ionization energy decreases moving down a group (increasing atomic radius).
The property is alternately still often called the ionization potential, measured in
kJ/mol .
For example, the first two molar ionization energies of magnesium (stripping the
two 3s electrons from a magnesium atom) are 738 and 1450 kJ/mol. The third
ionization energy is a much larger (7730 kJ/mol)
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Electron binding energy (BE) :
is the energy required to release an electron from its atomic or molecular
orbital when adsorbed to a surface rather than a free atom.
Binding energy values are normally reported as positive values with units of (ev).
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Electron affinity:
The ability of an atom to accept an electron. It is the energy change that
occurs when an electron is added to a gaseous atom.
Atoms with stronger effective nuclear charge have greater electron affinity.
The Group IIA elements, the alkaline earths, have low electron affinity
values(why)?
These elements are relatively stable because they have filled s subshells.
Group VIIA elements, the halogens, have high electron affinities because the
addition of an electron to an atom results in a completely filled shell.
Group VIII elements, noble gases, have electron affinities near zero, since each
atom possesses a stable octet and will not accept an electron readily.
Elements of other groups have low electron affinities.
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Electronegativity :
Electronegativity is a measure of the attraction of an atom for the electrons in a
chemical bond.
The higher the electronegativity of an atom, the greater its attraction for bonding
electrons. Electronegativity is related to ionization energy. Electrons with low
ionization energies have low electronegativities because their nuclei do not exert
a strong attractive force on electrons.
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Elements with high ionization energies have high electronegativities due to the
strong pull exerted on electrons by the nucleus.
Electronegativity increases on passing from left to right along a period,
In a group, the electronegativity decreases (as atomic number increases), as a
result of increased distance between the valence electron and nucleus (greater
atomic radius). An example of an electropositive (i.e., low electronegativity)
element is cesium; an example of a highly electronegative element is fluorine.
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