Arrange each of the following sets of atoms and ions in order of increasing size. Se2−,Te2−, Se

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Identify the periodic trend: Atomic and ionic size increases as you move down a group in the periodic table.

Recognize that Se and Te are in the same group (Group 16) with Te below Se, so Te is larger than Se.

Consider the effect of gaining electrons: Anions are larger than their neutral atoms because the added electrons increase electron-electron repulsion.

Compare the sizes: Se is smaller than Se²⁻ because Se²⁻ has gained electrons, increasing its size.

Arrange the atoms and ions in order of increasing size: Se < Se²⁻ < Te²⁻.

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Atomic Radius

The atomic radius is a measure of the size of an atom, typically defined as the distance from the nucleus to the outermost electron shell. Atomic size generally increases down a group in the periodic table due to the addition of electron shells, while it decreases across a period from left to right due to increased nuclear charge, which pulls electrons closer to the nucleus.

Ionic Radius

The ionic radius refers to the size of an ion in a crystal lattice. Cations (positively charged ions) are smaller than their parent atoms due to the loss of electron shells and increased effective nuclear charge, while anions (negatively charged ions) are larger due to the addition of electrons, which increases electron-electron repulsion in the outer shell.

Comparison of Ions and Atoms

When comparing the sizes of atoms and ions, it is essential to consider their charge and position in the periodic table. For example, Se2− and Te2− are both anions, with Te2− being larger due to its position below Se in the periodic table. In contrast, neutral Se is smaller than both anions, as it has fewer electrons and less electron-electron repulsion.

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Textbook Question

Consider S, Cl, and K and their most common ions. (a) List the atoms in order of increasing size.

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Textbook Question

Consider S, Cl, and K and their most common ions. (b) List the ions in order of increasing size. (c) Explain any differences in the orders of the atomic and ionic sizes.

Textbook Question

True or false? c. S2− is larger than K+.

Textbook Question

In the ionic compounds LiF, NaCl, KBr, and RbI, the measured cation–anion distances are 2.01 Å (Li–F), 2.82 Å (Na–Cl), 3.30 Å (K–Br), and 3.67 Å (Rb–I), respectively. b. Calculate the difference between the experimentally measured ion–ion distances and the ones predicted from Figure 7.8.

Textbook Question

In the ionic compounds LiF, NaCl, KBr, and RbI, the measured cation–anion distances are 2.01 Å (Li–F), 2.82 Å (Na–Cl), 3.30 Å (K–Br), and 3.67 Å (Rb–I), respectively. c. What estimates of the cation–anion distance would you obtain for these four compounds using neutral atom bonding atomic radii? Are these estimates as accurate as the estimates using ionic radii?