Arrange the atoms according to decreasing effective nuclear charge experienced by their valence electrons: S, Mg, Al, Si.

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Understand that effective nuclear charge \( (Z_{eff}) \) is the net positive charge experienced by valence electrons. It is calculated as \( Z_{eff} = Z - S \), where \( Z \) is the atomic number and \( S \) is the shielding constant.

Identify the atomic numbers: Sulfur (S) is 16, Magnesium (Mg) is 12, Aluminum (Al) is 13, and Silicon (Si) is 14.

Consider the electron configuration for each element to determine the shielding effect. The more inner electrons, the greater the shielding.

Recognize that elements in the same period have increasing \( Z_{eff} \) from left to right due to increasing atomic number and similar shielding.

Arrange the elements in order of decreasing \( Z_{eff} \): Sulfur (S), Silicon (Si), Aluminum (Al), Magnesium (Mg).

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주요 개념

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Effective Nuclear Charge (Z_eff)

Effective nuclear charge is the net positive charge experienced by an electron in a multi-electron atom. It accounts for the actual nuclear charge (number of protons) minus the shielding effect of inner-shell electrons. This concept is crucial for understanding how strongly valence electrons are attracted to the nucleus, influencing atomic size and reactivity.

Shielding Effect

The shielding effect refers to the phenomenon where inner-shell electrons partially block the attraction between the nucleus and the valence electrons. This results in a lower effective nuclear charge felt by the outermost electrons. Understanding this effect is essential for predicting trends in atomic properties across the periodic table.

Periodic Trends

Periodic trends are patterns observed in the properties of elements as you move across or down the periodic table. Key trends include atomic radius, ionization energy, and electronegativity, which are influenced by effective nuclear charge and shielding. Recognizing these trends helps in arranging elements based on their atomic characteristics, such as effective nuclear charge.