Arrange the following atoms in order of increasing effective nuclear charge experienced by the electrons in the n = 2 shell: Be, Br, Na, P, Se.

검증된 단계별 안내

Step 1: Understand that the effective nuclear charge is the net positive charge experienced by an electron in a multi-electron atom. The effective nuclear charge can be calculated using Slater's rules, but for this problem, we will use the concept of atomic number and electron shielding to determine the order.

Step 2: The n = 2 shell refers to the second energy level of an atom, which includes the 2s and 2p sublevels. For atoms in the same period of the periodic table, the effective nuclear charge generally increases from left to right. This is because as we move across a period, the number of protons (atomic number) increases, but the shielding effect remains relatively constant, leading to a higher effective nuclear charge.

Step 3: Be, Na, and P are in the second period of the periodic table, while Br and Se are in the fourth period. Within the second period, Be has the lowest atomic number, followed by Na and then P. Therefore, Be will have the lowest effective nuclear charge, followed by Na and then P.

Step 4: For atoms in the fourth period, Br and Se, the effective nuclear charge will be higher than those in the second period due to the higher atomic number. However, within the fourth period, Br has a lower atomic number than Se, so Br will have a lower effective nuclear charge than Se.

Step 5: Combining the information from steps 3 and 4, the atoms in order of increasing effective nuclear charge experienced by the electrons in the n = 2 shell are: Be, Na, P, Br, Se.

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주요 개념

질문에 올바르게 답하기 위해 반드시 이해해야 하는 핵심 개념들은 다음과 같습니다.

Effective Nuclear Charge (Z_eff)

Effective nuclear charge (Z_eff) is the net positive charge experienced by an electron in an atom, accounting for both the total nuclear charge and the shielding effect of other electrons. It can be calculated using the formula Z_eff = Z - S, where Z is the atomic number and S is the shielding constant. Higher Z_eff values indicate a stronger attraction between the nucleus and the electrons, influencing atomic size and ionization energy.

Shielding Effect

The shielding effect occurs when inner-shell electrons partially block the attraction between the nucleus and the outer-shell electrons. This results in a lower effective nuclear charge felt by the outer electrons. As the number of inner electrons increases, the shielding effect becomes more significant, leading to variations in atomic properties such as size and reactivity across the periodic table.

Trends in the Periodic Table

Periodic trends refer to predictable patterns in elemental properties as one moves across a period or down a group in the periodic table. For instance, effective nuclear charge generally increases across a period due to increasing atomic number and constant shielding, while it decreases down a group due to increased shielding from additional electron shells. Understanding these trends is essential for predicting the behavior of elements, including their effective nuclear charge.