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Chapter 2: Problem 9

With regard to electron configuration, what do all the elements in Group VIIA of the periodic table have in common?

Short Answer

Expert verified
Answer: The common feature in the electron configurations of Group VIIA elements is that they all have seven electrons in their outermost (valence) energy level with a "p^5" configuration. This explains their similar chemical properties, as they tend to gain one electron to achieve a stable electron configuration (a full octet).

Step by step solution

01

Identify the elements in Group VIIA

Group VIIA elements are: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), and Astatine (At).
02

Write the electron configurations for each element

To understand the commonality in Group VIIA elements, let's write down their electron configurations: - Fluorine (F): 1s^2 2s^2 2p^5 - Chlorine (Cl): 1s^2 2s^2 2p^6 3s^2 3p^5 - Bromine (Br): 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^5 - Iodine (I): 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6 5s^2 4d^10 5p^5 - Astatine (At): 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6 5s^2 4d^10 5p^6 6s^2 4f^14 5d^10 6p^5
03

Identify the common feature in electron configurations

By examining the electron configurations of Group VIIA elements, we can see that they all have seven electrons in their outermost (valence) energy level. The electron configuration of their valence shell ends with "p^5": - Fluorine (F): 2p^5 - Chlorine (Cl): 3p^5 - Bromine (Br): 4p^5 - Iodine (I): 5p^5 - Astatine (At): 6p^5 This "p^5" configuration is the common feature that all elements in Group VIIA share, which explains their similar chemical properties as they tend to gain one electron to achieve a stable electron configuration (a full octet).

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Key Concepts

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

Electron Configuration
The concept of "electron configuration" refers to the arrangement of electrons within an atom's electron shells and subshells. Each element has a unique electron configuration that helps dictate its chemical behavior and properties.

This configuration is written in a specific notation. It begins with the shell number (represented by the principal quantum number "n"), followed by the type of subshell (s, p, d, f), and ends with the number of electrons in that subshell indicated as a superscript.

For example, fluorine has the electron configuration: \(1s^2 2s^2 2p^5\). This tells us that it has 2 electrons in the 1s subshell, 2 electrons in the 2s subshell, and 5 electrons in the 2p subshell. Each electron configuration reveals vital information about the atom's valence electrons and its reactivity.
Valence Electrons
Valence electrons are the electrons present in the outermost shell of an atom. These are crucial because they are involved in chemical bonding and reactions.

Valence electrons determine an element's reactivity and how it interacts with other elements. In the case of Group VIIA elements, such as fluorine and chlorine, they have 7 valence electrons.

This is because their outer shell ends with a "p^5" configuration (e.g., \(2p^5\) for fluorine), which indicates they are one electron short of having a full octet. As a result, these elements frequently seek to gain an additional electron during chemical reactions to achieve stability, similar to the noble gases, which have a complete octet.
Periodic Table Group
The periodic table is organized into columns known as "groups," where elements with similar chemical and physical properties are aligned. Group VIIA, also known as Group 17, consists of halogens, which include fluorine, chlorine, bromine, iodine, and astatine.

Their placement in the periodic table signifies commonalities in their chemical behavior. This occurs primarily because they all have the same number of valence electrons. The elements in Group VIIA are highly reactive non-metals, known for forming salts when they react with metals.

Moreover, as you move down the group, the size of the atom increases, and their reactivity decreases. However, their electron affinity, or the tendency to gain electrons, remains a significant characteristic shared within the group.

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Most popular questions from this chapter

Without consulting Figure 2.6 or Table \(2.2,\) determine whether each of the electron configurations given below is an inert gas, a halogen, an alkali metal, an alkaline earth metal, or a transition metal. Justify your choices. (a) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{5}\) (b) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 3 d^{7} 4 s^{2}\) (c) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 3 d^{10} 4 s^{2} 4 p^{6}\) (d) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 4 s^{1}\) (e) $1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 3 d^{10} 4 s^{2} 4 p^{6} 4 d^{5} 5 s^{2}$ (f) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2}\)

Chromium has four naturally occurring isotopes: \(4.34 \%\) of \({ }^{50} \mathrm{Cr}\), with an atomic weight of \(49.9460\) amu; \(83.79 \%\) of \({ }^{52} \mathrm{Cr}\), with an atomic weight of \(51.9405 \mathrm{amu} ; 9.50 \%\) of \({ }^{53} \mathrm{Cr}\), with an atomic weight of \(52.9407 \mathrm{amu} ;\) and \(2.37 \%\) of \({ }^{54} \mathrm{Cr}\), with an atomic weight of \(53.9389\) amu. On the basis of these data, confirm that the average atomic weight of \(\mathrm{Cr}\) is \(51.9963 \mathrm{amu}\)

Explain why hydrogen fluoride (HF) has a higher boiling temperature than hydrogen chloride (HCl) \(\left(19.4^{\circ} \mathrm{C}\right.\) vs. \(\left.-85^{\circ} \mathrm{C}\right)\), even though HF has a lower molecular weight.

Sodium chloride \((\mathrm{NaCl})\) exhibits predominantly ionic bonding. The \(\mathrm{Na}^{+}\)and \(\mathrm{Cl}^{-}\)ions have electron structures that are identical to which two inert gases?

(a) Briefly cite the main differences between ionic, covalent, and metallic bonding. (b) State the Pauli exclusion principle.
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