Question

Hydrogen is an unusual element because it behaves in some ways like the alkali metal elements and in other ways like nonmetals. Its properties can be explained in part by its electron configuration and by the values for its ionization energy and electron affinity, (a) Explain why the electron affinity of hydrogen is much closer to the values for the alkali elements than for the halogens. (b) Is the following statement true? "Hydrogen has the smallest bonding atomic radius of any element that forms chemical compounds." If not, correct it. If it is, explain in terms of electron configurations. (c) Explain why the ionization energy of hydrogen is closer to the values for the halogens than for the alkali metals. (d) The hydride ion is \(\mathrm{H}\). Write out the process corresponding to the first ionization energy of hydride. (e) How does the process you wrote in part (d) compare to the process for the electron affinity of elemental hydrogen?

Short answer

(a) Hydrogen's electron affinity is closer to alkali metals because it has a similar electron configuration with one electron in its outermost shell. Gaining one electron stabilizes its electron configuration similar to noble gases. (b) Hydrogen has the smallest bonding atomic radius due to having only one electron in its 1s orbital which is closest to the nucleus. (c) Hydrogen's ionization energy is closer to halogens because they both have high ionization energies due to being close to achieving a stable electron configuration (complete outer shell). (d) The first ionization energy of hydride is represented as: \[\mathrm{H}^- (g) \rightarrow \mathrm{H} (g) + e^-\] (e) The electron affinity of elemental hydrogen and the ionization energy of hydride processes are essentially the reverse of each other, with the former releasing energy upon electron addition and the latter requiring energy for electron removal.

Step-by-step solution

(a) Electron affinity of hydrogen

The electron affinity of an element refers to the amount of energy released when an electron is added to a neutral atom in the gas phase to form a negative ion. Hydrogen's electron affinity is closer to alkali metals because it also belongs to Group 1 in the periodic table and shares a similar electron configuration. Hydrogen and alkali metals have only one electron in their outermost shell, and when they gain one electron, they achieve a stable electron configuration, similar to noble gases.

(b) Hydrogen's bonding atomic radius

The statement "Hydrogen has the smallest bonding atomic radius of any element that forms chemical compounds" is true. This is because hydrogen has only one electron in its 1s orbital, which is closest to the nucleus. Due to its small size and low atomic number, hydrogen has a smaller atomic radius than any other element.

(c) Ionization energy of hydrogen

Ionization energy is the energy required to remove an electron from an atom in its gaseous state. Hydrogen's ionization energy is closer to the values for the halogens because they both have high ionization energies. This is because hydrogen and halogens are closer to achieving a stable electron configuration (complete outer shell), and hence, it is more difficult to remove an electron from them compared to alkali metals, which have relatively low ionization energies.

(d) First ionization energy of hydride

The first ionization energy refers to the energy required to remove the first electron from a negatively charged ion. To represent the first ionization energy of hydride, we can write the following equation: \[\mathrm{H}^- (g) \rightarrow \mathrm{H} (g) + e^-\]

(e) Comparison to the electron affinity of elemental hydrogen

The process for the electron affinity of elemental hydrogen is the addition of an electron to create a hydride ion, which can be represented as: \[\mathrm{H} (g) + e^- \rightarrow \mathrm{H}^- (g)\] Comparing this process to the first ionization energy of hydride, we see that they are essentially the reverse of each other. The electron affinity of elemental hydrogen represents the energy released upon the addition of an electron, whereas the ionization energy of hydride represents the energy required to remove that same electron from the hydride ion.

Key concepts

Electron Affinity

Electron affinity is a measure of the energy change when an electron is added to a neutral atom in the gas phase, resulting in the formation of a negative ion. This process is generally exothermic, meaning it releases energy.
For hydrogen, its electron affinity is notably similar to that of alkali metals. This is because both hydrogen and alkali metals belong to Group 1 of the periodic table and exhibit similar simple electron configurations.

• Both have one electron in their outermost shell.
• Adding an electron leads them to achieve a noble gas-like electron configuration, which is more stable and energetically favorable.

In contrast, halogens have a much higher electron affinity since they are just one electron short of achieving a full valence shell, making them more eager to gain an electron compared to hydrogen.

Ionization Energy

Ionization energy refers to the energy required to remove an electron from a neutral atom in the gas phase. For hydrogen, this value is distinctively closer to that of the halogens rather than alkali metals.
This occurs because both hydrogen and halogens have relatively high ionization energies, meaning more energy is needed to remove an electron. These higher energies reflect their closer approach to a complete valence electron shell:

• Hydrogen has a single electron, and removing it implies breaking the balance involving its sole available electron.
• Halogens are one electron away from a full valence shell, making them naturally resistant to losing electrons.

Alkali metals, on the other hand, require much less energy for electron removal, due to their single electron being far from a full shell, leading to a lesser ionization energy.

Bonding Atomic Radius

The bonding atomic radius is the measure of the size of an atom that forms part of a single covalent bond. Hydrogen stands out in this respect, having the smallest bonding atomic radius of elements forming chemical compounds.
The primary reason is its simplest electron configuration—with just one electron occupying the 1s orbital, which is closest to the nucleus. This small size is due to:

• Its low atomic number, being the first element of the periodic table.
• Having no inner electron shell that would otherwise push the outer shell further out.

As such, when participating in bonding, hydrogen's electron clouds are tightly drawn to the positively charged nucleus, resulting in its notably small atomic radius.

Periodic Table

The periodic table is an organized chart showing all known chemical elements in a way that visually represents periodic trends. Hydrogen, despite being the simplest element with only one proton and one electron, is placed in Group 1 alongside alkali metals. This positioning can be intriguing because:

• Like alkali metals, hydrogen has a single electron in its outer shell, aligning it functionally rather than physically in the table.
• Hydrogen can also share properties with Group 17 (halogens), especially in its high ionization energy and electronegativity.

The periodic table layout allows observers to predict element behaviors based on position and correspondences, providing useful insights into element similarities and chemical behavior patterns.