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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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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
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:
- 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.
- 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
- 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 ).
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.
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)
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.
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.
