Choose the element with the higher first ionization energy from each pair. a. P or I b. Si or Cl c. P or Sb d. Ga or Ge

검증된 단계별 안내

Ionization energy is the energy required to remove an electron from an atom in the gaseous state.

Generally, ionization energy increases across a period (left to right) and decreases down a group (top to bottom) in the periodic table.

a. Compare P (Phosphorus) and I (Iodine): P is above I in the periodic table, so P has a higher ionization energy.

b. Compare Si (Silicon) and Cl (Chlorine): Cl is to the right of Si in the same period, so Cl has a higher ionization energy.

c. Compare P (Phosphorus) and Sb (Antimony): P is above Sb in the same group, so P has a higher ionization energy.

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Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. It is a key indicator of how strongly an atom holds onto its electrons. Generally, ionization energy increases across a period in the periodic table due to increasing nuclear charge and decreases down a group due to increased distance from the nucleus and electron shielding.

Periodic Trends

Periodic trends refer to the predictable patterns observed in the properties of elements as you move across or down the periodic table. For ionization energy, elements on the right side of the table (like Cl) typically have higher ionization energies than those on the left (like Si), while elements higher up in a group have higher ionization energies than those lower down (like P compared to Sb).

Comparison of Elements

When comparing elements for their ionization energies, it is essential to consider their positions in the periodic table. Elements in the same group exhibit similar properties, but the one higher up will usually have a higher ionization energy. Additionally, comparing elements from different groups requires understanding their electronic configurations and the effective nuclear charge experienced by their outermost electrons.