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Chapter 7: Problem 57

How are metallic character and first ionization energy related?

Short Answer

Expert verified
Metallic character and first ionization energy are inversely related due to underlying factors such as atomic size, shielding effect, and effective nuclear charge. As metallic character increases, ionization energy decreases, meaning elements with high metallic character can easily lose electrons, whereas elements with low metallic character have high ionization energies, making it more difficult for them to lose electrons.

Step by step solution

01

Define Metallic Character

Metallic character refers to the degree to which an element behaves like a metal. Metals tend to be ductile, malleable, and shine, and good conductors of electricity and heat. Their ability to lose electrons easily is a key aspect of their metallic character, as this explains many of their physical properties, as well as chemical reactivity.
02

Define First Ionization Energy

First ionization energy is the amount of energy required to remove the outermost valence electron from an isolated neutral atom in its gaseous state, forming a positive ion. Ionization energy is a measure of the strength of an atom's hold on its valence electron. It is dependent on factors such as atomic size, shielding effect, and effective nuclear charge.
03

Periodic Trends in Metallic Character and Ionization Energy

As we move across a period (from left to right), the metallic character of elements decreases, and the ionization energy increases. This is because the number of protons in the nucleus also increases, resulting in a stronger attraction to the electrons and a greater difficulty in their removal. As we move down a group, the metallic character increases and ionization energy decreases due to the increase in atomic size and the shielding effect of inner electrons. This results in an overall weakening of the attraction between the nucleus and the valence electrons.
04

Relationship between Metallic Character and First Ionization Energy

The metallic character and first ionization energy of an element are inversely related. Elements with high metallic character possess low ionization energies, which means they can easily lose their valence electrons. Similarly, elements with low metallic character have high ionization energies, making it more difficult for them to lose electrons to form positively charged ions. The relationship between these properties can be attributed to the influence of atomic size, shielding effect, and effective nuclear charge on both metallic character and ionization energy. In conclusion, metallic character and first ionization energy are inversely related due to the same underlying factors, such as atomic size, shielding effect, and effective nuclear charge, that influence both properties. Elements with high metallic character tend to have low ionization energies, while elements with low metallic character tend to have high ionization energies.

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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 the amount of energy required to strip away the outermost electron from a neutral atom in its gaseous form. It's a crucial measure of how strongly an atom holds onto its electrons.
An atom with high ionization energy holds its valence electron(s) tightly, making it challenging to remove them. In contrast, an atom with low ionization energy easily loses its electrons. Several factors influence the ionization energy, these include:
  • Atomic Size: Smaller atoms have higher ionization energies because the electrons are closer to the nucleus and more strongly attracted.
  • Shielding Effect: The presence of inner electrons can shield the outer electrons from the nucleus, reducing ionization energy.
  • Effective Nuclear Charge: As the number of protons in the atom increases, the pull on the electrons gets stronger, increasing ionization energy.
Periodic Trends
Understanding periodic trends is essential when studying the properties of elements, such as metallic character and ionization energy. These trends are predictable changes observable across the periodic table and can help infer properties of various elements.
As you move from left to right across a period, the atomic number increases, nuclear charge rises, and ionization energy generally increases. Simultaneously, metallic character decreases in these elements. Conversely, moving down a group follows a pattern where the atomic size increases, which results in a decreased ionization energy and increased metallic character.
This is because the growing number of energy levels and increased shielding effect cause the outer electrons to feel less attraction to the nucleus, making elements more metallic as you go down the group.
Atomic Size
Atomic size, or atomic radius, is the distance from the nucleus of an atom to the outer boundary of the surrounding cloud of electrons. It is a fundamental property used to explain many chemical behaviors.
Generally, atomic size decreases across a period from left to right. This happens because the increase in the number of protons leads to a greater nuclear charge, pulling the electron cloud closer. However, as we travel down a group, atomic size increases. New electron shells are added, outweighing the increase in nuclear charge.
Larger atomic size typically results in lower ionization energy. This is due to the electrons being further from the nucleus, making them easier to remove. Consequently, elements with a larger atomic radius often exhibit increased metallic character.
Shielding Effect
The shielding effect is a phenomenon where inner electron shells act as a barrier between the nucleus and the valence electrons. This effect causes a reduction in the nucleus' electrostatic pull on the outer electrons.
As more inner shells are added (moving down a group), the shielding effect increases. It counterbalances the increased nuclear charge and results in lower ionization energy for the outer electrons.
Conversely, across a period, the shielding effect remains relatively constant, while nuclear charge increases. This combination results in stronger attractions to the nucleus, higher ionization energy, and decreased metallic character. Understanding this concept helps explain why elements behave differently as you move across and down the periodic table.

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Most popular questions from this chapter

Does the reactivity of a metal correlate with its first ionization energy? Explain.

One way to measure ionization energies is ultraviolet photoelectron spectroscopy (UPS, or just PES), a technique based on the photoelectric effect. coo (Section 6.2 ) In PES, monochromatic light is directed onto a sample, causing electrons to be emitted. The kinetic energy of the emitted electrons is measured. The difference between the energy of the photons and the kinetic energy of the electrons corresponds to the energy needed to remove the electrons (that is, the ionization energy). Suppose that a PES experiment is performed in which mercury vapor is irradiated with ultraviolet light of wavelength \(58.4 \mathrm{nm}\). (a) What is the energy of a photon of this light in eV? (b) Write an equation that shows the process corresponding to the first ionization energy of \(\mathrm{Hg}\). (c) The kinetic energy of the emitted electrons is measured to be \(10.75 \mathrm{eV}\). What is the first ionization energy of Hg in kJ/mol? (d) Using Figure 7.9 , determine which of the halogen elements has a first ionization energy closest to that of mercury.

Until the early 1960 s the group 8 A elements were called the inert gases; before that they were called the rare gases. The term rare gases was dropped after it was discovered that argon accounts for roughly \(1 \%\) of Earth's atmosphere. (a) Why was the term inert gases dropped? (b) What discovery triggered this change in name? (c) What name is applied to the group now?

Use electron configurations to explain the following observations: (a) The first ionization energy of phosphorus is greater than that of sulfur. (b) The electron affinity of nitrogen is lower (less negative) than those of both carbon and oxygen. (c) The second ionization energy of oxygen is greater than the first ionization energy of fluorine. (d) The third ionization energy of manganese is greater than those of both chromium and iron.

(a) As described in Section \(7.7,\) the alkali metals react with hydrogen to form hydrides and react with halogens- for example, fluorine to form halides, Compare the roles of hydrogen and the halogen in these reactions. How are the forms of hydrogen and halogen in the products alike? (b) Write balanced equations for the reaction of fluorine with calcium and for the reaction of hydrogen with calcium. What are the similarities among the products of these reactions?
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