List the following ions in order of increasing radius Li^+,...

Key Concepts

Ionic Radius

Ionic radius refers to the effective distance from the nucleus to the outermost shell of electrons in an ion. Its size is influenced by factors such as the number of electron shells, effective nuclear charge, and electron-electron repulsions. Understanding ionic radius is crucial for predicting and comparing the sizes of different ions.

Effective Nuclear Charge

Effective nuclear charge is the net positive charge experienced by an electron in an atom or ion after accounting for the shielding effects of other electrons. A higher effective nuclear charge pulls electrons closer to the nucleus, resulting in a smaller ionic radius, while a lower effective nuclear charge allows electrons to be more spread out, increasing the size of the ion.

Electron Shielding and Principal Quantum Number

Electron shielding occurs when inner electrons reduce the full attractive force of the nucleus on the outer electrons. The principal quantum number indicates the energy level (and average distance from the nucleus) of these electrons. Increased shielding and higher principal quantum numbers generally lead to larger ionic radii because the outer electrons experience a weaker attraction to the nucleus.

Cation Versus Anion Size

Ions can be classified as cations (positively charged) or anions (negatively charged), and their sizes vary accordingly. Cations, which have lost electrons, typically have a smaller radius due to a higher effective nuclear charge acting on fewer electrons. Anions, having gained electrons, exhibit a larger radius because the added electron-electron repulsion and increased shielding allow the electron cloud to expand.

Transcript

00:01 So for this problem, we're asked to compare the atomic radii of the four ions i have written here to the right.

00:07 Before we do that, i just want to reiterate the general trends of atomic radius as they pertain to the periodic table.

00:15 So in general, atomic radius is going to increase from top to bottom as you move down a group, and it's going to increase from right to left as you move across a period.

00:27 So the two main reasons for this have to do with the size of the atom as well as the number of protons that it has.

00:34 So first, the size of the atom, as you consider looking at a group, as you move down a group, the n number for the atom is increasing.

00:44 So the valent shell is existing in increasingly higher energy orbitals, which are further from the nucleus.

00:53 So that's one reason, one factor that's driving the atomic radius larger.

00:58 The other thing to consider is the number of protons in the nucleus.

01:01 So if you look over to the far right side of the periodic table, these atoms have a higher number of protons relative to their counterparts within the same period to the left.

01:16 And more protons overall is going to further condense down those atomic radii and make them tighter because they're going to pull those electrons in just a little bit harder.

01:27 So keeping those trends in mind, we want to take a look at the ions that we have written here to the right and try and think about which ones are going to be the largest and which ones are going to be the smallest.

01:39 So in general, the charges here, i think, are meant to be a little bit of a deterring factor.

01:45 They're really actually not going to play a huge role in what we're considering here because the trends are actually going to follow the pretty general trends that we have here drawn to the left...

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