Which of the following statements about effective nuclear charge for the outermost valence electron of an atom is incorrect? (i) The effective nuclear charge can be thought of as the true nuclear charge minus a screening constant due to the other electrons in the atom. (ii) Effective nuclear charge increases going left to right across a row of the periodic table. (iii) Valence electrons screen the nuclear charge more effectively than do core electrons. (iv) The effective nuclear charge shows a sudden decrease when we go from the end of one row to the beginning of the next row of the periodic table. (v) The change in effective nuclear charge going down a column of the periodic table is generally less than that going across a row of the periodic table

Verified step by step guidance

1. Step 1: Understand the concept of effective nuclear charge. It is the net positive charge experienced by an electron in a multi-electron atom. The effective nuclear charge may be approximated by the equation Zeff = Z - S, where Z is the atomic number and S is the screening constant, roughly equivalent to the number of inner electrons.
2. Step 2: Analyze statement (i). This statement is correct. The effective nuclear charge can indeed be thought of as the true nuclear charge (the number of protons in the nucleus) minus a screening constant due to the other electrons in the atom.
3. Step 3: Analyze statement (ii). This statement is also correct. As we move from left to right across a period in the periodic table, the number of protons increases, but the number of inner (core) electrons remains the same. Therefore, the effective nuclear charge experienced by the valence electrons increases.
4. Step 4: Analyze statement (iii). This statement is incorrect. Core electrons are closer to the nucleus and shield the valence electrons from the nuclear charge more effectively than the valence electrons can shield each other. Therefore, core electrons screen the nuclear charge more effectively than do valence electrons.
5. Step 5: Analyze statements (iv) and (v). Both of these statements are correct. The effective nuclear charge shows a sudden decrease when we go from the end of one row to the beginning of the next row of the periodic table because the number of energy levels (and thus shielding electrons) increases. The change in effective nuclear charge going down a column of the periodic table is generally less than that going across a row of the periodic table because the number of protons increases at a slower rate.

Key Concepts

Effective Nuclear Charge (Z_eff)

Effective nuclear charge (Z_eff) is the net positive charge experienced by an electron in a multi-electron atom. It is calculated as the actual nuclear charge (Z) minus the shielding or screening effect caused by other electrons. This concept helps explain trends in atomic size and ionization energy across the periodic table.

Shielding Effect

The shielding effect refers to the phenomenon where inner electrons partially block the attraction between the nucleus and the outermost (valence) electrons. This results in a lower effective nuclear charge felt by the valence electrons compared to the actual nuclear charge. Understanding this effect is crucial for analyzing how atomic properties change across periods and groups.

Periodic Trends

Periodic trends are patterns observed in the properties of elements as one moves across a period or down a group in the periodic table. Key trends include atomic radius, ionization energy, and electronegativity, which are influenced by effective nuclear charge and electron shielding. Recognizing these trends is essential for predicting the behavior of elements in chemical reactions.