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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Ch.7 - Periodic Properties of the Elements
Chapter 7, Problem 35
Provide a brief explanation for each of the following: (a) Cl- is larger than Ar. (b) P3- is larger than S2-. (c) K+ is larger than Na+. (d) F- is larger than F.
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Consider the position of K and Na in the periodic table. K (potassium) is located in period 4, while Na (sodium) is in period 3.
Understand that as you move down a group in the periodic table, the number of electron shells increases. This means K has more electron shells than Na.
Recognize that both K+ and Na+ ions have lost one electron compared to their neutral atoms, resulting in the same electron configuration as the noble gas preceding them.
Despite having the same electron configuration, K+ has more electron shells than Na+, making it larger in size.
Conclude that the increase in the number of electron shells as you move down a group results in a larger ionic radius for K+ compared to Na+.
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Key Concepts
Here are the essential concepts you must grasp in order to answer the question correctly.
Ionic Radius
Ionic radius refers to the size of an ion in a crystal lattice. Cations, like K<sup>+</sup> and Na<sup>+</sup>, are formed when atoms lose electrons, resulting in a decrease in size due to reduced electron-electron repulsion and increased effective nuclear charge. The ionic radius can vary based on the ion's charge and the number of electron shells.
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Periodic Trends
Periodic trends are patterns observed in the properties of elements across the periodic table. As you move down a group, the ionic radius increases due to the addition of electron shells, which outweighs the increase in nuclear charge. This explains why K<sup>+</sup>, located below Na<sup>+</sup> in Group 1, is larger.
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Effective Nuclear Charge
Effective nuclear charge (Z<sub>eff</sub>) is the net positive charge experienced by an electron in a multi-electron atom. It accounts for the shielding effect of inner electrons. In the case of K<sup>+</sup> and Na<sup>+</sup>, although both are cations, K<sup>+</sup> has more electron shells, leading to a larger ionic radius despite a similar Z<sub>eff</sub>.
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Related Practice
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
Arrange each of the following sets of atoms and ions, in orderof increasing size: (a) Pb, Pb2+, Pb4+
Textbook Question
In the ionic compounds LiF, NaCl, KBr, and RbI, the measured cation–anion distances are 201 pm (Li–F), 282 pm (Na–Cl), 330 pm (K–Br), and 367 pm (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 201 pm (Li–F), 282 pm (Na–Cl), 330 pm (K–Br), and 367 pm (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?
Textbook Question
Write equations that show the processes that describe thefirst, second, and third ionization energies of a chlorineatom. Which process would require the least amount ofenergy?