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

Based on their positions in the periodic table, predict which atom of the following pairs will have the smaller first ionization energy: (a) $\mathrm{Br}, \mathrm{Kr} ; \mathbf{( b )} \mathrm{C}, \mathrm{Ca} ;(\mathbf{c}) \mathrm{Li}, \mathrm{Rb} ;$; (d) $\mathrm{Pb}, \mathrm{Si} ;$ (e) $\mathrm{Al}, \mathrm{B}$.

Based on their positions in the periodic table, predict which atom of the following pairs will have the smaller first ionization energy:
(a) $\mathrm{Br}, \mathrm{Kr} ; \mathbf{( b )} \mathrm{C}, \mathrm{Ca} ;(\mathbf{c}) \mathrm{Li}, \mathrm{Rb} ;$; (d) $\mathrm{Pb}, \mathrm{Si} ;$
(e) $\mathrm{Al}, \mathrm{B}$.

Key Concepts

Atomic Radius

Atomic radius is the measure of the size of an atom, typically defined as the distance from the nucleus to the outermost electron shell. Generally, atoms with larger radii have lower ionization energies because their outer electrons are farther from the nucleus and more shielded, making them easier to remove.

Ionization Energy

Ionization energy is the energy required to remove an electron from a neutral atom in its gaseous state. It is a fundamental property that reflects how strongly an atom holds onto its electrons and is influenced by the atom's nuclear charge, electron configuration, and electron-electron repulsions.

Periodic Trends

Periodic trends are patterns in the properties of elements that emerge when elements are arranged by increasing atomic number. Trends such as increasing ionization energy across a period and decreasing ionization energy down a group occur due to changes in effective nuclear charge and atomic size.

Effective Nuclear Charge

Effective nuclear charge refers to the net positive charge experienced by an electron in a multi-electron atom. It is calculated by subtracting the effects of electron shielding from the full nuclear charge, and it increases across a period, leading to higher ionization energies.

Electron Shielding

Electron shielding is the phenomenon where inner-shell electrons reduce the full attractive force of the nucleus on the outer electrons. This effect becomes more pronounced in larger atoms and results in lower effective nuclear charge felt by the outermost electrons, thereby reducing ionization energy.

Transcript

00:01 This question is going to take a look at determining an important periodic trend of ionization energy.

00:08 So ionization energy is the energy required to remove an electron from a substance.

00:16 And so we're going to do some comparisons of ionization energy through some different sets of atoms or elements.

00:26 The first set we're going to look at is comparing bromine and krypton.

00:33 General rule that we want to follow for ionization energy is as we move left to right across the periodic table, we see an increase in ionization energy.

00:44 And as you move top to bottom down any family on the periodic table, we see a decrease in ionization energy.

00:53 So we'll kind of keep that on the back burner as we do some of these comparisons, including this very first one.

00:59 When we compare bromine to krypton, we see that krypton.

01:03 Is to the right of bromine on the periodic table.

01:07 And that means it should have an higher ionization energy than bromine.

01:12 So if you're to put these in order of ionization energy, krypton then would have the higher ionization energy than bromine according to our trend.

01:24 The same thing with another set of examples we're going to compare here, silicon to let.

01:33 That's not what i meant to write.

01:35 Let's try that again here...

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