Log out Go to App
Go to App

Learning Materials

Features

Discover

Chapter 19: Problem 4

Use the shell model of atoms and the concept of core charge to explain why the AVEE of atoms generally decreases from top to bottom down a group.

Short Answer

Expert verified
The Atomic Vertical Electron Energy (AVEE) tends to decrease down a group due to the increase in the number of electron shells and the corresponding increase in distance between the valence electrons and the nucleus. This results in lesser attraction between the two, thus reducing the AVEE. Additionally, inner shell electrons shield the valence electrons from the full effect of the core charge, further reducing the AVEE.

Step by step solution

01

Understanding the Concept of AVEE and Core Charge

The total energy of an atom's electrons (the atomic vertical electron energy or AVEE) is dependent on two factors: the attraction between the electrons and the nucleus (which depends on the charge of the nucleus, or the core charge), and the repulsion between the electrons themselves. As one moves down the groups in the periodic table, the number of electron shells increases, effectively increasing the distance between the nucleus and the valence electrons.
02

Linking the Shell Model of Atoms to AVEE

According to the shell model of atoms, electrons are arranged in shells around the nucleus. Each succeeding shell is located farther from the nucleus. Therefore, as we go from top to bottom in a group, the number of shells increases, and so does the distance between the valence electrons and the nucleus.
03

Explaining the decrease in AVEE

The increased distance between the valence electrons and the nucleus (due to more shells) means that the attraction between the electrons and the nucleus decreases. Meanwhile, electrons in inner shells shield the valence electrons from the full effect of the core charge. Consequently, the AVEE decreases as we go from top to bottom down a group in the periodic table. That means electrons at higher levels (toward the bottom of a group) generally have less energy associated to them compared to those at lower levels (towards the top).

Unlock Step-by-Step Solutions & Ace Your Exams!

  • Full Textbook Solutions

    Get detailed explanations and key concepts

  • Unlimited Al creation

    Al flashcards, explanations, exams and more...

  • Ads-free access

    To over 500 millions flashcards

  • Money-back guarantee

    We refund you if you fail your exam.

Start your free trial

Over 30 million students worldwide already upgrade their learning with Vaia!

One App. One Place for Learning.

All the tools & learning materials you need for study success - in one app.

Get started for free

Most popular questions from this chapter

Use the shell model of atoms and the concept of core charge to explain why the AVEE of atoms generally increases from left to right across a period.

Describe the trend in EN moving left to right across a period of the periodic table.

In terms of AVEE, explain why Na is more likely to form \(\mathrm{Na}^{+}\) than \(\mathrm{O}\) is to form \(\mathrm{O}^{+}\)

a) There is one valence electron in a \(\mathrm{H}\) atom. What is the average ionization energy of the one valence electron in \(\mathrm{H}\) ? (Note that this is a very easy question because there is only one valence electron!) b) There are two valence electrons in a He atom. What is the average ionization energy of the two valence electrons in He? (Again, an easy question because both valence electrons have the same IE!) c) Which atom, \(\mathrm{H}\) or He, holds its valence electron(s) "more tightly" on average? Explain your reasoning.

Describe the trend in EN moving down a group of the periodic table.
See all solutions

Recommended explanations on Chemistry Textbooks

Organic Chemistry

Read Explanation

Nuclear Chemistry

Read Explanation

Chemical Analysis

Read Explanation

Chemical Reactions

Read Explanation

Physical Chemistry

Read Explanation

Inorganic Chemistry

Read Explanation
View all explanations

What do you think about this solution?

We value your feedback to improve our textbook solutions.

Study anywhere. Anytime. Across all devices.

Sign-up for free