$$U = \text{energy required to break up an ionic solid into individual gaseous ions}$$

Question 1

Which of the following statements best explains why cations are smaller than their corresponding neutral atoms?

Accepted Answer

Cations have fewer electrons, reducing electron–electron repulsion and often resulting in a lower principal quantum number for the valence shell.

Question 2

Which of the following ions does NOT have a true noble-gas electron configuration, but instead has an additional filled d subshell?

Accepted Answer

Zn$$^{2+}$$

Question 3

Which of the following best describes the trend in ionic radii for anions compared to their neutral atoms?

Accepted Answer

Anions are larger because of additional electron–electron repulsions and a decrease in effective nuclear charge.

Question 4

Which equation correctly represents the first ionization energy for sodium?

Accepted Answer

$$\mathrm{Na(g)} \rightarrow \mathrm{Na$$^{+}$$(g)} + \mathrm{e$$^{-}$$}$$

Question 5

Which group of elements has slightly larger first ionization energies than expected, and why?

Accepted Answer

Group 2A (Be, Mg, Ca), because their filled s subshells are more stable and harder to remove electrons from.

Question 6

Which of the following statements about electron affinity is correct?

Accepted Answer

A negative electron affinity means energy is released when an electron is added.

Question 7

Which element would you expect to have a zero electron affinity, and why?

Accepted Answer

Neon, because it has a filled valence shell and does not readily accept electrons.

Question 8

According to the octet rule, which of the following is most likely to form an ion with a noble-gas configuration?

Accepted Answer

Potassium, by losing one electron to form $$\mathrm{K$$^{+}$$}$$

Question 9

Which of the following best describes the Born-Haber cycle for the formation of an ionic solid?

Accepted Answer

It is a series of steps that calculates the lattice energy by considering ionization energy, electron affinity, and other enthalpy changes.

Question 10

Which equation represents the lattice energy ($$U) $$for an ionic solid?

Accepted Answer

$$U = 	ext{energy required to break up an ionic solid into individual gaseous ions}$$

Question 11

A student is asked to explain why metals tend to lose electrons in chemical reactions. Which reasoning is most accurate?

Accepted Answer

Metals have low ionization energies and low electron affinities, making it easier for them to lose electrons.

Question 12

Which of the following would result in a large jump in ionization energy for a main-group element?

Accepted Answer

Removing an electron from a filled s or p subshell after the valence electrons have been removed.

Question 13

Which of the following best explains why group 6A elements (O, S) have slightly smaller first ionization energies than expected?

Accepted Answer

The first electron is removed from a filled orbital, which is less stable due to electron–electron repulsion.

Question 14

Which of the following is a real-world application of lattice energy?

Accepted Answer

Predicting the solubility of table salt ($$\mathrm{NaCl}) in $$water based on its lattice energy.

Question 15

Which of the following statements about the octet rule is correct?

Accepted Answer

Main-group elements tend to react to achieve eight outer-shell electrons, resulting in filled s and p sublevels.

Question 16

A chemist is analyzing the formation of $$\mathrm{MgO}$$ using the Born-Haber cycle. Which step involves the electron affinity of oxygen?

Accepted Answer

Adding an electron to $$\mathrm{O(g)} to $$form $$\mathrm{O$$^{-}$$(g)}$$

Question 17

Which of the following best describes the electron configuration of $$\mathrm{Na$$^{+}$$}$$?

Accepted Answer

$$1s^2$$\ $$2s^2$$\ $$2p^6$$

Question 18

Which of the following scenarios would result in the largest lattice energy?

Accepted Answer

A compound formed from small, highly charged ions such as $$\mathrm{Mg$$^{2+}$$}$$ and $$\mathrm{O$$^{2-}$$}$$

Ionic Compounds: Periodic Trends and Bonding Theory

Study Guide - Practice Questions