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Chapter 8: Problem 10

What is penetration? How is it related to shielding? Use the penetration effect to explain the difference in relative orbital energies of a \(3 p\) and a \(3 d\) electron in the same atom.

Short Answer

Expert verified
Penetration is the ability of an electron to approach the nucleus, closely related to shielding. Due to greater penetration, 3p electrons are lower in energy than 3d electrons in the same atom.

Step by step solution

01

- Define Penetration

Penetration in chemistry refers to the ability of an electron in an atom to get close to the nucleus. Electrons in orbitals with higher penetration experience a greater nuclear charge.
02

- Define Shielding

Shielding refers to the reduction of the effective nuclear charge on an electron due to the presence of other electrons between it and the nucleus. Electrons in inner shells shield outer electrons from the full charge of the nucleus.
03

- Relationship Between Penetration and Shielding

Penetration and shielding are interconnected. An electron that penetrates closer to the nucleus is less shielded by other electrons and thus experiencing a greater effective nuclear charge.
04

- Compare Penetration of 3p and 3d Electrons

The 3p orbital has more penetration than the 3d orbital due to its shape and spatial distribution. This means electrons in the 3p orbital can get closer to the nucleus compared to electrons in the 3d orbital.
05

- Explain the Relative Orbital Energies

Because 3p electrons penetrate more, they experience a higher effective nuclear charge and are thus lower in energy compared to 3d electrons. As a result, in the same atom, a 3p orbital is lower in energy than a 3d orbital.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

electronic penetration
Electronic penetration describes the ability of an electron to approach the atomic nucleus. Significant penetration means the electron is getting very close to the nucleus. This matters because electrons near the nucleus experience a strong attraction from the protons in the nucleus. For example, in a 2s orbital, the electron can get closer to the nucleus compared to an electron in a 2p orbital. This is due to the shape and spatial distribution of the orbital. The closer the electron, the more electrical attraction it feels.
effective nuclear charge
Effective nuclear charge is the net positive charge experienced by an electron in an atom. This is different from the actual nuclear charge because electrons are not only attracted by the protons but also repelled by other electrons. To calculate it, we take the total number of protons and subtract the screening effects of inner, shielding electrons. So, if we have an atom with a nuclear charge of +6 and 2 shielding electrons, the effective nuclear charge would be +4. More penetration by an electron results in a higher effective nuclear charge.
orbital energies
Orbital energies describe the energy levels of electrons in different orbitals. Remarkably, the energy depends on both the shape of the orbital and the effective nuclear charge. For instance, a 2s electron will generally have lower energy than a 2p electron because it penetrates the nucleus more effectively and experiences a higher effective nuclear charge. The same trend continues with higher orbitals, where the shape and penetration abilities dictate the relative energies.
shielding effect
The shielding effect arises because electrons in lower energy levels (inner electrons) protect outer electrons from the full effect of the nucleus's positive charge. Imagine a crowded room where people nearer to the speaker (nucleus) can hear clearly, but those farther away (outer electrons) hear less clearly because of the interference (shielding) from people in between. This is why electrons further from the nucleus feel a weaker effective nuclear charge. Shielding makes it easier for these outer electrons to escape since they don't feel the full attraction of the nucleus.
3p and 3d orbitals
Let's compare 3p and 3d orbitals. Both are in the same principal energy level ( = 3), but they have different shapes and penetration abilities. The 3p orbital penetrates the nucleus more than the 3d orbital. As a result, a 3p electron feels a higher effective nuclear charge and is lower in energy compared to a 3d electron. This explains why when filling an atom's orbitals, the 3p orbital gets filled before the 3d orbital. The 3p electrons are closer to the nucleus and hence, more stable due to higher effective nuclear charge.

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Most popular questions from this chapter

How many electrons in an atom can have each of the following quantum numbers or sublevel designations? (a) \(n=2, l=1\) (b) \(3 d\) (c) \(4 s\)

Name the element described in each of the following: (a) Smallest atomic radius in Group \(6 \mathrm{~A}(16)\) (b) Largest atomic radius in Period 6 (c) Smallest metal in Period 3 (d) Highest IE \(_{1}\) in Group \(4 \mathrm{~A}(14)\) (e) Lowest IE \(_{1}\) in Period 5 (f) Most metallic in Group \(5 \mathrm{~A}(15)\) (g) Group \(3 \mathrm{~A}(13)\) element that forms the most basic oxide (h) Period 4 element with highest energy level filled (i) Condensed ground-state electron configuration of [Ne] \(3 s^{2} 3 p^{2}\) (j) Condensed ground-state electron configuration of \([\mathrm{Kr}] 5 s^{2} 4 d^{6}\) (k) Forms \(2+\) ion with electron configuration [Ar] \(3 d^{3}\) (1) Period 5 element that forms \(3+\) ion with pseudo-noble gas configuration (m) Period 4 transition element that forms \(3+\) diamagnetic ion (n) Period 4 transition element that forms \(2+\) ion with a half-filled \(d\) sublevel (o) Heaviest lanthanide (p) Period 3 element whose \(2-\) ion is isoelectronic with Ar (q) Alkaline earth metal whose cation is isoelectronic with \(\mathrm{Kr}\) (r) Group \(5 \mathrm{~A}(15)\) metalloid with the most acidic oxide

What is the key distinction between sublevel energies in oneelectron species, such as the \(\mathrm{H}\) atom, and those in many-electron species, such as the \(\mathrm{C}\) atom? What factors lead to this distinction? Would you expect the pattern of sublevel energies in \(\mathrm{Be}^{3+}\) to be more like that in \(\mathrm{H}\) or that in \(\mathrm{C}\) ? Explain.

Write the full electron configuration of the Period 2 element with the following successive IEs (in \(\mathrm{kJ} / \mathrm{mol}\) ): \(\mathrm{IE}_{1}=801\) \(\mathrm{IE}_{2}=2427\) $$ \mathrm{IE}_{3}=3659 $$ \(\mathrm{IE}_{4}=25,022 \quad \mathrm{IE}_{5}=32,822\)

Which element would you expect to be less metallic? (a) Cs or Rn (b) Sn or Te (c) Se or Ge
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