Using data from Appendix 6, estimate bond lengths for the...

Using data from Appendix $6,$ estimate bond lengths for the gas phase diatomics $\mathrm{F}_{2}, \mathrm{Cl}_{2}, \mathrm{Br}_{2}$ and $\mathrm{I}_{2}$ Compare your answers with the values listed in Table $4.10 .$ Suggest reasons for the trend in bond lengths.

Key Concepts

Bond Length

Bond length is the distance between the nuclei of two bonded atoms. It is a fundamental concept in molecular structure and can often be estimated by summing the covalent radii of the atoms involved in the bond, providing insight into the size of the molecule.

Covalent Radius

The covalent radius represents the effective size of an atom when it is bonded to another atom. It is typically derived from experimental measurements and is used to estimate bond lengths by approximately representing the half-distance between two bonded nuclei.

Periodic Trends in Atomic Size

Periodic trends, especially the increase in atomic size down a group, are essential in predicting bond lengths. As atoms increase in size due to added electron shells, the distance between their nuclei when bonded tends to increase, leading to longer bond lengths for heavier elements.

Electron-Electron Repulsion

Electron-electron repulsion, notably from non-bonding (lone pair) electrons, plays a significant role in influencing bond lengths. Increased repulsion can cause bonds to stretch slightly beyond the sum of the ideal covalent radii, affecting the overall bond distance and stability.

Orbital Overlap Efficiency

Effective orbital overlap is key in forming strong chemical bonds. In larger atoms, the orbitals are more diffuse and may not overlap as efficiently, which can result in longer bond lengths. This concept helps explain trends in bond lengths as atomic size increases.

Transcript

00:01 Okay, so this question is just asking us about, we have diatomic molecules.

00:09 They look something like this, and we're trying to find the length of the bond.

00:16 That's all we're doing, and we're trying to compare these answers to table 410.

00:22 So if i go up here, i have the name, the radius, given in a couple different things from appendix 6, the bond length, and table 410, which we'll just used to check.

00:36 So the first thing we need to do is figure out which radius of a single atom do we do we use.

00:45 We have the vanderwall, covalent, metallic, or ionic.

00:49 It wouldn't make any sense to use metallic because they're gaseous.

00:54 It wouldn't make any sense to use ionic because they don't have a charge.

00:59 So we're stuck with vanderwallin and covalent.

01:02 But since these are bonds and not single atoms, we don't want to use vanderwalls.

01:07 We want to use our covalent radii.

01:12 So we go to table, uh, we go to appendix 6 and we find the covalent radius of each of these.

01:20 And so for fluorine, it's 71 picometers.

01:29 For chlorine, it's 99 picometers.

01:35 For bromine, it's 114 and for iodine it's 133.

01:45 Okay? so these are just two atoms per molecule.