Based on their positions in the periodic table, predict which has the smallest first ionization

Based on their positions in the periodic table, predict...

Key Concepts

Atomic Radius and Electron Shielding

Atomic radius and electron shielding significantly influence ionization energy. Larger atomic radii usually mean that valence electrons are farther from the nucleus and less tightly held due to increased shielding, which results in lower ionization energy.

Effective Nuclear Charge

Effective nuclear charge is the net positive charge experienced by the electrons after accounting for the shielding effect of inner-shell electrons. A higher effective nuclear charge means that electrons are held more tightly, leading to higher ionization energies.

Ionization Energy

Ionization energy refers to the amount of energy required to remove an electron from an atom in its gaseous state. This concept is fundamental to understanding how atoms interact chemically and how easily they can form cations.

Periodic Trends

Periodic trends describe the predictable changes in properties such as ionization energy across periods and groups in the periodic table. Generally, ionization energy increases across a period and decreases down a group due to changes in both effective nuclear charge and electron shielding.

Transcript

00:01 So for this problem, we're given a list of atoms, and we want to identify the one with the smallest ionization energy.

00:06 To do that, we should first consider the general trends of ionization energy within the periodic table.

00:11 And those are that ionization energy is going to increase as you move from bottom to top within a group.

00:16 That has a lot to do with the fact that the atoms at the bottom of a group have a larger atomic radius, which means that those electrons in the valence shell are held pretty far away from the nucleus, and their attraction to the positive charge of the nucleus is pretty weak.

00:31 So it's much easier to pull off a 7 -s electron from francium than it is to pull off a 6s -cesium electron.

00:41 That same principle is actually applicable to the other trend of the periodic table where ionization energy increases from left to right.

00:49 So by the same rule, the size of the atom plays somewhat of a role, but another thing to consider with these, is the fact that they have a larger number of protons condensed within the same energy level.

01:06 So we look at the second period here, for example, all of these valence electrons are held within the n -equal -2 energy level, and the number of protons are increasing from left to right.

01:17 So because fluorine has nine protons condensed in its nucleus, all holding onto an n -equal -2 energy level of electrons, it's much more difficult to pull an electron from the outermost shell of fluorine than it is from the outermost shell of nitrogen, for example.

01:36 And again, that has to do with that increasing proton count...

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