MCAT General Chemistry Review - Alexander Stone Macnow, MD 2019-2020

The Periodic Table
Explanations to Discrete Practice Questions

1. BThe periodic table is organized into periods (rows) and groups (columns). Groups (columns) are particularly significant because they represent sets of elements with the same valence electron configuration, which in turn will dictate many of the chemical properties of those elements. Although (A) is true, the fact that both ions are positively charged does not explain the similarity in chemical properties; most metals produce positively charged ions. (C) is not true because lithium and sodium are in the same group, not period. Finally, although lithium and sodium do have relatively low atomic weights, so do several other elements that do not share the same properties, eliminating (D).

2. AAs one moves from top to bottom in a group (column), extra electron shells accumulate, despite the fact that the valence configurations remain identical. These extra electron shells provide shielding between the positive nucleus and the outermost electrons, decreasing the electrostatic attraction and increasing the atomic radius. Because carbon and silicon are in the same group, and silicon is farther down in the group, silicon will have a larger atomic radius because of its extra electron shell.

3. CAtomic radius is determined by multiple factors. Of the choices given, the number of valence electrons does have an impact on the atomic radius. As one moves across a period (row), protons and valence electrons are added, and the electrons are more strongly attracted to the central protons. This attraction tightens the atom, shrinking the atomic radius. The number of electron shells is also significant, as demonstrated by the trend when moving down a group (column). As more electron shells are added that separate the positively charged nucleus from the outermost electrons, the electrostatic forces are weakened, and the atomic radius increases. The number of neutrons is irrelevant because it does not impact these attractive forces.

4. CIonization energy increases from left to right, so the first ionization energy of lithium is lower than that of beryllium. Second ionization energy is always larger than first ionization energy, so beryllium’s second ionization energy should be the highest value. This is because removing an additional electron from Be+ requires one to overcome a significantly larger electrostatic force.

5. BAntimony (Sb) is on the right side of the periodic table, but not far right enough to be a nonmetal, (D). It certainly does not lie far enough to the right to fall in Group VIIA (Group 17), which would classify it as a halogen, (C). While sources have rarely classified antimony as a metal, (A), it is usually classified as a metalloid, (B).

6. CElectronegativity describes how strong an attraction an element will have for electrons in a bond. A nucleus with a larger effective nuclear charge will have a higher electronegativity; Zeff increases toward the right side of a period. A stronger nuclear pull will also lead to increased first ionization energy, as the forces make it more difficult to remove an electron. The vertical arrow can be explained by the size of the atoms. As size decreases, the positive charge becomes more effective at attracting electrons in a chemical bond (higher electronegativity), and the energy required to remove an electron (ionization energy) increases.

7. CAll four descriptions of metals are true, but the most significant property that contributes to the ability of metals to conduct electricity is the fact that they have valence electrons that can move freely. Malleability, (A), is the ability to shape a material with a hammer, which does not play a role in conducting electricity. The low electronegativity and high melting points of metals, (B) and (D), also do not play a major role in the conduction of electricity.

8. BThis block represents the alkaline earth metals, which form divalent cations, or ions with a +2 charge. All of the elements in Group IIA have two electrons in their outermost s subshell. Because loss of these two electrons would leave a full octet as the outermost shell, becoming a divalent cation is a stable configuration for all of the alkaline earth metals. Although some of these elements might be great conductors, they are not as effective as the alkali metals, eliminating (A). (C) is also incorrect because, although forming a divalent cation is a stable configuration for the alkaline earth metals, the second ionization energy is still always higher than the first. Finally, (D) is incorrect because atomic radii increase when moving down a group of elements because the number of electron shells increases.

9. BIron is a transition metal. Transition metals can often form more than one ion. Iron, for example, can be Fe2+ or Fe3+. The transition metals, in these various oxidation states, can often form hydration complexes with water. Part of the significance of these complexes is that, when a transition metal can form a complex, its solubility within the related solvent will increase. The other ions given might dissolve readily in water, but because none of them are transition metals, they will not likely form complexes.

10. DThis question is simple if one recalls that periods refer to the rows in the periodic table, while groups or families refer to the columns. Within the same period, an additional valence electron is added with each step toward the right side of the table.

11. BThis question requires knowledge of the trends of electronegativity within the periodic table. Electronegativity increases as one moves from left to right for the same reasons that effective nuclear charge increases. Electronegativity decreases as one moves down the periodic table because there are more electron shells separating the nucleus from the outermost electrons. In this question, chlorine is the furthest toward the top-right corner of the periodic table.

12. BElectron affinity is related to several factors, including atomic size and filling of the valence shell. As atomic radius increases, the distance between the nucleus and the outermost electrons increases, thereby decreasing the attractive forces between protons and electrons. As a result, increased atomic radius will lead to lower electron affinity. Because atoms are in a low-energy state when their outermost valence electron shell is filled, atoms needing only one or two electrons to complete this shell will have high electron affinities. In this example, (B) and (D) need only one more electron to have a noble gas-like electron configuration; because (B) is smaller, it will have the highest electron affinity.

13. DThe effective nuclear charge refers to the strength with which the protons in the nucleus can pull on electrons. This phenomenon helps to explain electron affinity, electronegativity, and ionization energy. In (A), the nonionized chlorine atom, the nuclear charge is balanced by the surrounding electrons: 17 p+/17 e. The chloride ion, (B), has a lower effective nuclear charge because there are more electrons than protons: 17 p+/18 e. Next, elemental potassium, (C), has the lowest effective nuclear charge because it contains additional inner shells that shield its valence electron from the nucleus. (D), ionic potassium, has a higher effective nuclear charge than any of the other options do because it has the same electron configuration as Cl (and the same amount of shielding from inner shell electrons as neutral Cl) but contains two extra protons in its nucleus: 19 p+/18 e.

14. DIonic bonds are formed through unequal sharing of electrons. These bonds typically occur because the electron affinities of the two bonded atoms differ greatly. For example, the halogens have high electron affinities because adding a single electron to their valence shells would create full valence shells. In contrast, the alkaline earth metals have very low electron affinities and are more likely to be electron donors because the loss of two electrons would leave them with full valence shells. (A) states the opposite and is incorrect because the halogens have high electron affinity and the alkaline earth metals have low electron affinity. (B) is incorrect because equal sharing of electrons is a classic description of covalent bonding, not ionic. (C) is a true statement, but is not relevant to why ionic bonds form.

15. CWhen n = 3, l = 0, 1, or 2. The highest value for l in this case is 2, which corresponds to the d subshell. Although the 3d block appears to be part of the fourth period, it still has the principal quantum number n = 3. In general, the subshells within an energy shell increase in energy as follows: s < p < d < f (although there is no 3f subshell).