Log out Go to App
Go to App

Learning Materials

Features

Discover

Chapter 7: Problem 42

(a) What is the trend in first ionization energies as one proceeds down the group 17 elements? Explain how this trend relates to the variation in atomic radii. (b) What is the trend in first ionization energies as one moves across the fourth period from \(\mathrm{K}\) to \(\mathrm{Kr}\) ? How does this trend compare with the trend in atomic radii?

Short Answer

Expert verified
\( (a) \) The first ionization energies decrease as one proceeds down the group 17 elements, as the atomic radii increase because outermost electrons are less strongly attracted to the nucleus. \( (b) \) As one moves across the fourth period from K to Kr, the first ionization energies increase, contrary to the trend in atomic radii which decrease. This is due to the stronger attraction between the nucleus and the electrons, as the number of protons increases.

Step by step solution

01

Ionization Energy and Atomic Radii

Ionization energy is the energy required to remove an electron from an atom (or ion) in the gaseous state. As for atomic radius, it is defined as the distance between the nuclei of two adjacent atoms. Step 2: Analyzing the Trend in First Ionization Energies in Group 17 elements
02

Trend in Group 17 elements

Group 17 elements are also known as halogens. As we move down the group, the atomic radii increase as there are more electron shells present in each successive element. With a larger atomic radius, the electrons in the outermost shell are further away from the nucleus and are less strongly attracted to it. Consequently, less energy is required to remove an electron from the outermost shell. Therefore, the first ionization energies decrease as one proceeds down the group 17 elements. Step 3: Relating the Trend in Ionization Energies with Atomic Radii for Group 17 elements
03

Relation between Ionization Energies and Atomic Radii in Group 17

The trend in first ionization energies in group 17 elements is directly related to the variation in atomic radii. As the atomic radii increase down the group, the first ionization energies decrease due to the outermost electrons being further away from the nucleus and being less strongly attracted to it. Step 4: Analyzing the Trend in First Ionization Energies across the Fourth Period
04

Trend in the Fourth Period

As we move across the fourth period from K (potassium) to Kr (krypton), the atomic number increases which implies that the number of protons in the nucleus increases. Consequently, the positive charge in the nucleus becomes stronger. This results in a stronger attraction between the nucleus and the electrons. This stronger attraction makes it more difficult to remove an electron, meaning the ionization energies increase from left to right across the fourth period. Step 5: Comparing the Trend in Ionization Energies with the Trend in Atomic Radii across the Fourth Period
05

Comparison in Fourth Period

The trend in first ionization energies across the fourth period is contrary to the trend in atomic radii. As we move from K to Kr, the atomic radii decrease because the increasing number of protons in the nucleus pulls the electrons closer to it. However, the ionization energies increase from left to right across the fourth period due to the stronger attraction between the nucleus and the electrons.

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

Give three examples of +2 ions that have an electron configuration of $n d^{10}(n=3,4,5 \ldots)$

(a) Why does xenon react with fluorine, whereas neon does not? (b) Using appropriate reference sources, look up the bond lengths of Xe-F bonds in several molecules. How do these numbers compare to the bond lengths calculated from the atomic radii of the elements?

(a) Why is calcium generally more reactive than beryllium? (b) Why is calcium generally less reactive than rubidium?

Moseley's experiments on \(\mathrm{X}\) rays emitted from atoms led to the concept of atomic numbers. (a) If arranged in order of increasing atomic mass, which element would come after chlorine? (b) Describe two ways in which the properties of this element differ from the other elements in group $8 \mathrm{~A}$.

Consider \(\mathrm{S}, \mathrm{Cl},\) and \(\mathrm{K}\) and their most common ions. (a) List the atoms in order of increasing size. (b) List the ions in order of increasing size. (c) Explain any differences in the orders of the atomic and ionic sizes.
See all solutions

Recommended explanations on Chemistry Textbooks

Nuclear Chemistry

Read Explanation

Inorganic Chemistry

Read Explanation

The Earths Atmosphere

Read Explanation

Physical Chemistry

Read Explanation

Chemical Reactions

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.

Sign-up for free