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Home , Core charge, Effective nuclear charge, Shielding effect

Effective Nuclear Charge :

The is the net positive charge experienced by an in a multi-electron . The term "effective" is used because the of negatively charged 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 . It is possible to determine the strength of the nuclear charge by looking at the oxidation number of the atom.

1 2 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 in the nucleus (), 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*. 3 Values Shielding effect :

The shielding effect describes the decrease in attraction between an electron and the nucleus in any atom with more than one . 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 4 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) . 5 Periodic properties : :

The atomic radius of an element is half of the distance between the centers of two 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.

6 Moving down a group in the , 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.

7 : 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 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) 8 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). 9 :

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.

10 : 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 because their nuclei do not exert a strong attractive force on electrons. 11 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 .

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