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Chapter 7: Problem 45

Which of the following has more acidic character? (a) \(\mathrm{H}_{2} \mathrm{Se}\) (b) \(\mathrm{H}_{2} \mathrm{~S}\) (c) \(\mathrm{H}_{2} \mathrm{O}\) (d) \(\mathrm{H}_{2} \mathrm{Te}\)

Short Answer

Expert verified
Hydrogen telluride (\text{H}_{2} \text{Te}) has the most acidic character among the given options.

Step by step solution

01

Consider the Periodic Trends for Acidity

The acidic character of hydrides increases with the size of the central atom within the same group of the periodic table. This is due to the increased ability to donate the hydrogen ion (H+) as the size of the central atom increases, thereby making the bond with hydrogen weaker and more polarizable.
02

Compare the Position of Elements in the Periodic Table

Looking at the elements provided, Selenium (Se), Sulfur (S), Oxygen (O), and Tellurium (Te), they are all in the same group of the periodic table (group 16). As you move down the group from Oxygen to Tellurium, the size of the atoms increases.
03

Determine the Acid with Greater Acidic Character

Since Te is the largest atom among Se, S, and O, within the same group, its corresponding hydride, H2Te, will have the weakest bond to hydrogen and therefore can more easily donate H+ ions, making it the most acidic.

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Key Concepts

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

Periodic Trends in Acidity
Understanding the periodic trends in acidity is vital in predicting the acidic character of different compounds, particularly hydrides like those in the given exercise. Generally, acidity in hydrides increases as we move down a group in the periodic table.

This occurs because the atomic size of the central atom becomes larger, weakening its bond with hydrogen. A weaker bond indicates that it’s easier for the compound to donate hydrogen ions, the essence of what it means to be an acid. For example, in group 16, the hydrides become more acidic from H2O to H2Te due to this trend. This concept highlights the importance of knowing an element's position in the periodic table to predict its properties.
Donating Hydrogen Ions
When discussing donating hydrogen ions, we're looking at an element's tendency to release H+ into solution—a key characteristic of acids. This ion donation is critical to understanding acidity.

The ability of a hydride to donate a hydrogen ion relates to bond strength; weaker bonds between hydrogen and the central atom mean ions can be donated more readily. In our exercise, compounds like H2Te have larger central atoms, leading to longer and weaker bonds, which facilitates easier hydrogen ion donation and thus greater acidity.
Size of Central Atom in Acidity
The size of the central atom plays a crucial role in determining the acidity of a hydride. The larger the central atom, the more polarizable the electron cloud, which in turn weakens the bond with hydrogen.

This weaker bond is less capable of holding onto the hydrogen ion, making the hydride a better proton donor or a stronger acid. In the context of our exercise, while all the elements belong to the same group and have similar electron configurations, their atomic size increases from oxygen to tellurium, making H2Te the most acidic hydride among the given options due to the large size of the tellurium atom.
Comparison of Elements in the Periodic Table
Effective comparison of elements in the periodic table is an indispensable skill for chemistry students. It allows us to deduce various properties and reactivity trends.

In our case, comparing group 16 elements—such as sulfur and selenium—shows that as atomic number increases, so does the size of the atom, which affects its chemical behavior. It's not just the family relation (group membership) that matters, but also the period or row placement. It is this relative placement in the periodic table that leads to the informed conclusion that among the choices, H2Te stands out as having the most acidic character.

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