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Chapter 3: Problem 77

What is the trend in first ionization energy as you go down a given group in the periodic table? As you go across a period from left to right?

Short Answer

Expert verified
As you move down a given group in the periodic table, the first ionization energy decreases due to larger atomic radius and increased electron shielding. Across a period from left to right, the first ionization energy increases because of the increased effective nuclear charge and smaller atomic radius, leading to stronger attraction between electrons and the nucleus.

Step by step solution

01

1. Understanding Ionization Energy

Ionization energy refers to the amount of energy needed to remove an electron from an atom or ion in its gaseous state. Generally, the ionization energies of elements depend on factors such as electron shielding, atomic radius, and effective nuclear charge.
02

2. First Ionization Energy Trend Down a Group

As you move down a group in the periodic table, the number of electron shells increases, and the atomic radius becomes larger. Consequently, the valence electrons (the outermost electrons) are further away from the nucleus and are not held as tightly to the nucleus because of increased shielding from inner electrons. Therefore, the first ionization energy typically decreases as you move down a group in the periodic table.
03

3. First Ionization Energy Trend Across a Period

When moving across a period from left to right, there is an increase in the effective nuclear charge as the number of protons in the nucleus increases. With a stronger effective nuclear charge, the outermost electrons feel a stronger attraction to the nucleus, making them more difficult to remove. Additionally, as we move across a period, the atomic radius decreases, further increasing the effective nuclear charge felt by the valence electrons. Consequently, the first ionization energy increases as you move across a period from left to right. In summary:
04

4. Summary of Ionization Energy Trends

- The trend in first ionization energy as you move down a given group in the periodic table is decreasing because the atomic radius increases, resulting in weaker attraction to the nucleus and increased electron shielding. - The trend in first ionization energy as you go across a period from left to right in the periodic table is increasing due to the increased effective nuclear charge and smaller atomic radius, causing stronger attraction between electrons and the nucleus.

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Key Concepts

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

First Ionization Energy
First ionization energy is a crucial concept in chemistry that indicates the amount of energy required to remove the outermost electron from a neutral atom in its gaseous state.

Understanding this energy helps us predict how likely an element is to participate in chemical reactions involving electron exchange. For instance, elements with low ionization energies readily lose electrons to form positive ions, making them reactive, particularly metals. Conversely, elements with high ionization energies, such as nonmetals, are less likely to part with their electrons, becoming negative ions in chemical reactions.
Periodic Table Groups
The periodic table groups are vertical columns that categorize elements with similar properties. As you observe the elements within a group, there are patterns in their reactivity, electron configuration, and atomic structure.

Each element in a group has the same number of valence electrons, which contributes to these shared chemical behaviors. Valence electrons play a pivotal role in bonding, so understanding group trends is essential. For instance, alkali metals of Group 1 have one valence electron, which they readily lose to react aggressively with others, especially water.
Effective Nuclear Charge
Effective nuclear charge is a term that refers to the net positive charge experienced by an electron in a multi-electron atom. The concept illustrates that the full nuclear charge is not felt by an electron due to the shielding effect from other electrons in inner shells.

As the number of protons in the nucleus (atomic number) increases, so does the effective nuclear charge, leading to a stronger attraction of the nucleus for the valence electrons. This increased attraction impacts the atom's properties, including ionization energy and atomic radius.
Atomic Radius
The atomic radius is the distance from the atom's nucleus to the boundary of its surrounding electron cloud. This radius plays a vital role in determining an atom's chemical properties, such as reactivity and the strength of bonds it can form.

Within the periodic table, atomic radii tend to decrease as you move from left to right across a period, due to the increase in effective nuclear charge attracting electrons closer to the nucleus. In contrast, atomic radii increase as you move down a group, since more electron shells are added, which takes electrons further from the nucleus.
Electron Shielding
Electron shielding is the phenomenon where inner electrons reduce the effective nuclear charge felt by outer electrons. It arises because the inner electrons act as a 'shield' between the nucleus and the valence electrons. This effect becomes more pronounced as the number of electron shells increases, impacting ionization energies.

The increased shielding in larger atoms means that outer electrons are not as strongly attracted to the nucleus, which typically makes them easier to remove. Thus, the concept of electron shielding helps explain why ionization energy decreases going down a group in the periodic table.

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Most popular questions from this chapter

How does a group differ from a period in the periodic table?

Suppose we wanted to use one of the more massive atoms, like \({ }_{92}^{235} \mathrm{U}\), as the standard reference for the atomic mass scale. (a) What would be the universally agreed-upon atomic mass of the \({ }^{235} \mathrm{U}\) isotope? (b) Why don't we use \({ }_{92}^{235} \mathrm{U}\) instead of \({ }_{6}^{12} \mathrm{C}\) ?

List three properties of metals and two properties of nonmetals. Describe the behavior of metalloids relative to the behaviors of metals and nonmetals.

Write the full atomic symbol for each isotope: (a) 8 protons, 9 neutrons (b) Atomic number 50 , mass number 119 (c) 12 neutrons, atomic number 11 (d) 28 protons, mass number 58 (e) 81 neutrons, symbol Ba

Explain why the first and second statements of Dalton's atomic theory are not exactly true.
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