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Chapter 7: Problem 89

Elements in group \(7 \mathrm{~A}\) in the periodic table are the halogens; elements in group \(6 \mathrm{~A}\) are called the chalcogens. (a) What is the most common oxidation state of the chalcogens compared to the halogens? Can you suggest an explanation for the difference? (b) For each of the following periodic properties, state whether the halogens or the chalcogens have larger values: atomic radii; ionic radii of the most common oxidation state; first ionization energy; second ionization energy.

Short Answer

Expert verified
(a) The most common oxidation state of chalcogens is -2, while for halogens, it is -1. The difference is due to the number of valence electrons in their outer shell, with chalcogens needing to gain two electrons and halogens needing one to achieve a stable noble gas electron configuration. (b) Chalcogens have larger atomic radii and ionic radii in their most common oxidation state (-2) compared to halogens. Halogens have larger first and second ionization energies than chalcogens due to the increasing effective nuclear charge across a period.

Step by step solution

01

Part (a): Identifying the most common oxidation states

Chalcogens are in group 6A, which means they have six valence electrons. To achieve a stable noble gas electron configuration, they can gain two electrons, giving them a common oxidation state of -2. Halogens are in group 7A, so they have seven valence electrons. To achieve a stable noble gas electron configuration, they just need to gain one electron, which results in a common oxidation state of -1.
02

Part (a): Explanation for the difference in oxidation states

The difference in the most common oxidation state between chalcogens and halogens boils down to the number of valence electrons in their outer shell. Chalcogens have six valence electrons, so they gain two electrons to have a full octet, achieving an oxidation state of -2. Halogens, on the other hand, have seven valence electrons and need only one more to reach the full octet, resulting in an oxidation state of -1.
03

Part (b): Comparing atomic radii

In general, atomic radii increase going down a group and decrease going across a period. Since chalcogens are to the left of halogens in the periodic table, they will generally have larger atomic radii compared to halogens.
04

Part (b): Comparing ionic radii of the most common oxidation state

When an atom gains electrons, it forms a negatively charged ion. In their common oxidation states, both chalcogens and halogens form negative ions, with chalcogens forming -2 ions and halogens forming -1 ions. In general, ions with more negative charges will have larger ionic radii due to increased electron repulsion. Therefore, the ionic radii of chalcogens in their most common oxidation state will be larger compared to that of halogens.
05

Part (b): Comparing first ionization energies

Ionization energy is the energy required to remove an electron from an atom. As you move across a period from left to right, ionization energy generally increases due to an increasing effective nuclear charge. Therefore, halogens will typically have larger first ionization energies compared to chalcogens.
06

Part (b): Comparing second ionization energies

Second ionization energy is the energy required to remove a second electron from an atom. Similar to the first ionization energy, the second ionization energy will be higher for halogens than chalcogens due to the increasing effective nuclear charge across a period.

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Most popular questions from this chapter

How do the sizes of atoms change as we move (a) from left to right across a row in the periodic table, (b) from top to bottom in a group in the periodic table? (c) Arrange the following atoms in order of increasing atomic radius: \(\mathrm{O}, \mathrm{Si}, \mathrm{I}, \mathrm{Ge} .\)

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.

We can draw an analogy between the attraction of an electron to a nucleus and seeing a lightbulb -in essence, the more \(n u=\) clear charge the electron "sees," the greater the attraction. (a) Within this analogy, discuss how the screening by core electrons is analogous to putting a frosted-glass lampshade between the lightbulb and your eyes, as shown in the illustration. (b) Explain how we could mimic moving to the right in a row of the periodic table by changing the wattage of the lightbulb. (c) How would you change the wattage of the bulb and/or the frosted glass to mimic the effect of moving down a column of the periodic table? [Section 7.2]

Compare the elements bromine and chlorine with respect to the following properties: (a) electron configuration, (b) most common ionic charge, (c) first ionization energy, (d) reactivity toward water, (e) electron affinity, (f) atomic radius. Account for the differences between the two elements.

Discussing this chapter, a classmate says, "Since elements that form cations are metals and elements that form anions are nonmetals, elements that do not form ions are metalloids." Do you agree or disagree? Explain your answer.
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