When comparing the size of different ions, the general radii...

Name: Problem 24, Chapter 3: Chemistry An Atoms First Approach
Uploaded: 2020-07-04T12:07:34-08:00
Duration: 6 min 59 s
Channel: Bryan Li
Description: When comparing the size of different ions, the general radii trend discussed in Chapter 2 is usually not very useful. What do you concentrate on when comparing sizes of ions to each other or when comparing the size of an ion to its neutral atom?

When comparing the size of different ions, the general radii trend discussed in Chapter 2 is usually not very useful. What do you concentrate on when comparing sizes of ions to each other or when comparing the size of an ion to its neutral atom?

Key Concepts

Effective Nuclear Charge

This concept refers to the net positive charge felt by electrons in an atom or ion after accounting for all the shielding effects from other electrons. When comparing ionic sizes, it is crucial to consider how losing or gaining electrons changes the effective nuclear charge, thereby influencing the attraction between the nucleus and the electron cloud. An ion with a higher effective nuclear charge tends to pull its electrons closer, resulting in a smaller radius compared to its neutral atom.

Electron Configuration and Shielding

Understanding the arrangement of electrons, including the impact of electron-electron repulsion and shielding, is key when comparing the sizes of ions or of an ion and its neutral atom. Changes in electron configuration upon ionization alter how the electrons shield each other, which in turn affects the size of the charge cloud. For instance, the removal of an electron in a cation decreases electron-electron repulsion and increases the effective nuclear charge experienced by the remaining electrons, leading to a contraction of the electron cloud.

Transcript

00:01 All right, in this question, we are comparing periodic trends of ions versus their native molecules, of their native atoms, sorry.

00:08 So we know that normally as you go across the periodic table, the radius of these atoms decrease.

00:16 And as you go down the periodic table, the radius increases.

00:20 Now, why is that? so let's look at, we'll do boron as an example.

00:28 Boron has an s -orbital, a 2 -s -orbital, and a p -orbital.

00:40 And a p -orbital doesn't look like a circle, a sphere.

00:43 It's a bunch of lobes, but it's hard to draw.

00:46 So i'm going to represent the p -orbital as a blue circle.

00:50 Now, in the s -orbital, boron has two atoms, i mean two electrons.

00:54 In the 2s orbital, boron has another two electrons, and in the p orbital, it's got one more electron.

01:07 So the electron configuration of boron is 1s2, 2s2, and 2p1.

01:21 Okay? and let's look at what nitrogen looks like.

01:31 Has 1s2, 2s2, and 2p3.

01:40 Because it's got two more electrons than the boron, and so it's going to go in the highest unoccupied molecular orbital.

01:52 So nitrogen has an s orbital, a 2s orbital, and then it's got its p orbital.

02:04 And if we draw in the electrons, we've got two here.

02:07 Here, two here, then one, two, three.

02:14 Okay.

02:16 Now, why is nitrogen smaller than boron? that's because of shielding.

02:25 So this boron has the two s orbitals shielding, or the p orbital.

02:37 And it's got five protons, which are pulling all of these electrons in.

02:49 Nitrogen has seven protons, and it's got the same amount of shielding.

03:00 So these protons on these electrons on the outer rim are going to be pulled stronger than the single electron and this boron, because nitrogen has more protons.

03:11 Okay, so that explains why as you go across, the radius goes down, and then as you go down, the radius increases because you have more orbitals, right? because if we look at aluminum, for example, it's not only got all these orbitals full, but it's got a 3s orbital and so on...

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