/*! This file is auto-generated */ .wp-block-button__link{color:#fff;background-color:#32373c;border-radius:9999px;box-shadow:none;text-decoration:none;padding:calc(.667em + 2px) calc(1.333em + 2px);font-size:1.125em}.wp-block-file__button{background:#32373c;color:#fff;text-decoration:none} Problem 76 Why do the metallic elements of ... [FREE SOLUTION] | 91影视

91影视

Log out

Learning Materials

Features

Discover

Chapter 10: Problem 76

Why do the metallic elements of a given period (horizontal row) typically have much lower ionization energies than do the nonmetallic elements of the same period?

Short Answer

Expert verified
In a given period, metallic elements have lower ionization energies than non-metallic elements due to their larger atomic sizes, greater shielding effect, and smaller effective nuclear charges. These factors result in a weaker attraction between the nucleus and the outermost electrons in metallic elements, making it easier to remove an electron and resulting in lower ionization energy.

Step by step solution

01

Understand Ionization Energy

Ionization energy is the energy required to remove an electron from a neutral atom in its gaseous state. It is dependent on the atomic size, shielding effect, and the effective nuclear charge experienced by the electron. Elements positioned on the left side of the periodic table (metals) typically have lower ionization energies, while elements on the right side (nonmetals) have higher ionization energies.
02

Atomic Size

As we move across a period from left to right, the atomic size decreases. This is because the number of protons in the nucleus increases, making the positive charge more concentrated which attracts the electrons closer to the nucleus. However, metallic elements have larger atomic sizes as compared to non-metallic elements in the same period. The larger atomic size makes it easier to remove an electron, resulting in lower ionization energy.
03

Shielding Effect

The shielding effect is the reduction in the effective nuclear charge experienced by the outermost electrons due to the presence of inner electrons. The shielding effect remains fairly constant across a period since there are no additional electron shells. However, the metallic elements have more shielding due to the presence of more inner electrons. This causes the outer electrons to feel a weaker pull from the nucleus, which contributes to lower ionization energy.
04

Effective Nuclear Charge

Effective nuclear charge (Z_eff) is the net positive charge experienced by an electron in an atom. As we move from left to right in a period, Z_eff increases because the number of protons in the nucleus increases. The metallic elements, which are on the left of the periodic table, experience smaller Z_eff compared to non-metallic elements. As a result, the outermost electrons in metallic elements are less attracted to the nucleus, giving them lower ionization energy.
05

Conclusion

In summary, metallic elements in a given period have lower ionization energies than non-metallic elements due to their larger atomic sizes, greater shielding effect, and smaller effective nuclear charges. These factors make it relatively easier to remove an electron from metallic elements compared to non-metallic elements.

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.

Over 30 million students worldwide already upgrade their learning with 91影视!

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

Although a hydrogen atom has only one electron, the hydrogen atom possesses a complete set of available orbitals. What purpose do these additional orbitals serve?

In each of the following sets of elements, indicate which element has the smallest atomic size. a. \(\mathrm{Ba}, \mathrm{Ca}, \mathrm{Ra}\) b. \(\mathrm{P}, \mathrm{Si}, \mathrm{Al}\) c. \(\mathrm{Rb}, \mathrm{Cs}, \mathrm{K}\)

Which of the following orbital designations is (are) not correct? a. \(1 d\) b. \(2 d\) c. \(4 d\) d. \(5 f\)

In each of the following sets of elements, which element would be expected to have the highest ion ization energy? a. \(\mathrm{Cs}, \mathrm{K}, \mathrm{Li}\) b. Ba, Sr, Ca c. I, Br, Cl d. \(\mathrm{Mg}, \mathrm{Si}, \mathrm{S}\)

The portion of the electromagnetic spectrum between wavelengths of approximately 400 and 700 nanometers is called the _______ region.
See all solutions

Recommended explanations on Chemistry Textbooks

Ionic and Molecular Compounds

Read Explanation

Kinetics

Read Explanation

Chemical Analysis

Read Explanation

Chemistry Branches

Read Explanation

Nuclear Chemistry

Read Explanation

Making Measurements

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.