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Chapter 7: Problem 44

Answer the following questions about the elements with the electron configurations shown here: $$ A=[A r] 4 s^{2} \quad B=[A r] 3 d^{10} 4 s^{2} 4 p^{5} $$ (a) Is element A a metal, metalloid, or nonmetal? (b) Is element \(B\) a metal, metalloid, or nonmetal? (c) Which element is expected to have the larger ionization energy? (d) Which element has the smaller atomic radius?

Short Answer

Expert verified
(a) Metal, (b) Nonmetal, (c) Element B, (d) Element B.

Step by step solution

01

Analyze Electron Configuration of Element A

The electron configuration for element A is \([Ar] 4s^2\). This shows that element A has two electrons in its outermost s-orbital after the argon core, indicating it is likely at the start of a new period on the periodic table with a filled 's' subshell and no 'p' electrons. Such a configuration is typical of alkaline earth metals, which are metals. Therefore, element A is a metal.
02

Analyze Electron Configuration of Element B

The electron configuration for element B is \([Ar] 3d^{10} 4s^2 4p^5\). This configuration indicates that element B is from a group near the end of a period with a nearly full 'p' subshell. The filled d-subshell also suggests it belongs to the p-block of the periodic table. This characteristic is typical of halogens, which are nonmetals. Therefore, element B is a nonmetal.
03

Compare Ionization Energies

Ionization energy tends to increase across a period from left to right. With element A being an alkaline earth metal and element B being a halogen, both in the same period, the element on the right (B) generally has a higher ionization energy than the element on the left (A). Therefore, element B is expected to have the larger ionization energy.
04

Compare Atomic Radii

Atomic radius decreases across a period from left to right due to increasing nuclear charge that pulls electrons closer to the nucleus. Since element A is an alkaline earth metal and element B is a halogen, element A is on the left side and element B is on the right side of the period. Therefore, element B is expected to have a smaller atomic radius than element A.

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Key Concepts

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

Periodic Table
The periodic table organizes all known elements in a meaningful way. It arranges the elements in order of increasing atomic number, which is the number of protons in the nucleus of an atom.
Each horizontal row is called a period, and each vertical column is called a group. Elements are often grouped based on similar properties.
  • Metals, found on the left and center, are generally shiny, good conductors of electricity and heat, and are malleable.
  • Nonmetals, located on the right side, are more varied in appearance and are typically insulators, meaning they do not conduct electricity well.
  • Metalloids, which include elements like silicon and germanium, have properties of both metals and nonmetals.
By knowing an element's position on the periodic table, we can infer a great deal about its electron configuration, reactive properties, and the types of compounds it forms.
Ionization Energy
Ionization energy is the energy required to remove an electron from an atom in the gas phase. It's an important concept because it helps explain an element's reactivity.
On the periodic table, ionization energy trends show that:
  • Ionization energy tends to increase as you move from left to right across a period. This is because the increased number of protons in the nucleus pulls the electrons closer, making them harder to remove.
  • Ionization energy tends to decrease as you move down a group. The electrons are further from the nucleus in higher energy levels, making them easier to remove.
For example, in the given exercise, element B, being a halogen, likely has a higher ionization energy than element A, an alkaline earth metal, because elements toward the end of a period generally exhibit higher ionization energies.
Atomic Radius
The atomic radius is half the distance between the nuclei of two atoms of the same element that are bonded together. It helps us understand an element's size.
In the periodic table:
  • Atomic radius decreases from left to right across a period. As elements gain more protons and electrons, the increased nuclear charge draws the electron cloud closer to the nucleus.
  • Atomic radius generally increases as you move down a group. Here, new electron shells are added, making the atoms larger despite the increasing nuclear charge.
Within the exercise, element A is likely to have a larger atomic radius than element B, since A is positioned toward the beginning of a period (alkaline earth metals) while B is near its end (halogens).
Metals and Nonmetals
Metals and nonmetals are differentiated by specific physical and chemical properties. Understanding these differences is key in chemistry:
  • Metals are typically solid at room temperature (except mercury), have high melting points, and can conduct electricity.
  • Nonmetals can be gases, liquids, or solids at room temperature, have lower melting points, and are poor conductors of electricity.
  • Metals tend to lose electrons in chemical reactions to form positive ions, while nonmetals tend to gain electrons to form negative ions.
In the exercise, element A is an alkaline earth metal, which means it possesses typical metallic properties. On the other hand, element B is a halogen, which is a group of nonmetals known for their high reactivity and tendency to gain electrons during reactions.

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Most popular questions from this chapter

Using orbital box diagrams and noble gas notation, depict the electron configurations of (a) \(\mathrm{V},\) (b) \(\mathrm{V}^{2+},\) and (c) \(\mathrm{V}^{5+} .\) Are any of the ions paramagnetic?

Using orbital box diagrams, depict an electron configuration for each of the following ions: (a) \(\mathrm{Na}^{+}\) (b) \(\mathrm{Al}^{3+},\) (c) \(\mathrm{Ge}^{2+},\) and \((\mathrm{d}) \mathrm{F}^{-}\).

Explain each answer briefly. (a) Place the following elements in order of increasing ionization energy: \(F, O,\) and \(S\) (b) Which has the largest ionization energy: O, S, or Se? (c) Which has the most negative electron attachment enthalpy: Se, Cl, or Br? (d) Which has the largest radius: \(\mathbf{O}^{2-}, \mathbf{F}^{-},\) or \(\mathrm{F} ?\)

Name the element corresponding to each characteristic below. (a) the element with the electron configuration \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{3}\) (b) the alkaline earth element with the smallest atomic radius (c) the element with the largest ionization energy in Group \(5 \mathrm{A}\) (d) the element whose \(2+\) ion has the configuration \([\mathrm{Kr}] 4 d^{5}\) (e) the element with the most negative electron attachment enthalpy in Group \(7 \mathrm{A}\) (f) the element whose electron configuration is \([\mathrm{Ar}] 3 d^{10} 4 s^{2}\)

The actinide americium, Am, is a radioactive element that has found use in home smoke detectors. Depict its electron configuration using noble gas and spdf notations.
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