C. The number decreases as the atomic number increases.
D. The number increases as the atomic number increases.
11. Which of the following elements has the highest electronegativity?
A. Mg
B. Cl
C. Zn
D. I
12. Of the four atoms depicted here, which has the highest electron affinity?
13. Which of the following is true of an atom with a large atomic radius?
A. The atom is most likely to be on the right side of the periodic table.
B. The atom is likely to have a high second ionization energy.
C. The atom is likely to have low electronegativity.
D. The atom is likely to form ionic bonds.
14. Which of the following atoms/ions has the largest effective nuclear charge?
A. Cl
B. Cl-
C. K
D. K+
15. Why do halogens often form ionic bonds with alkaline earth metals?
A. The alkaline earth metals have much higher electron affinities than the halogens.
B. By sharing electrons equally, the alkaline earth metals and halogens both form full octets.
C. Within the same row, the halogens have smaller atomic radii than the alkaline earth metals.
D. The halogens have much higher electron affinities than the alkaline earth metals.
16. What is the outermost orbital of elements in the third period?
A. s-orbital
B. p-orbital
C. d-orbital
D. f-orbital
Small Group Questions
1. Mercury (Hg) exists as a liquid at room temperature. Why, then, is it classified as a metal? What metallic properties might it possess?
2. The transition metals in group VIIIB could theoretically have eight different oxidation states. In reality, that does not hold true. Why not?
Explanations to Practice Questions
1. B
To answer this question, one must first recall that the periodic table is organized with periods (rows) and groups (columns). This method of organization allows elements to be organized such that some chemical properties can be predicted based on an element’s position in the table. Groups (columns) are particularly significant because they represent sets of elements with the same outer electron configuration. In other words, all elements within the same group will have the same configuration of valence electrons, which in turn will dictate many of the chemical properties of those similar elements. Although (A) is true, the fact that both ions are positively charged does not explain the similarity in chemical properties as effectively as does answer (B); most other metals, whether or not they are similar to lithium and sodium, produce positively charged ions. (C) is not true, because periods are rows and lithium and sodium are in the same column. Finally, although lithium and sodium do have relatively low atomic weights, so do several other elements that do not share those same properties.
2. A
This question assesses understanding of a key periodic trend: atomic radii. As one moves from left to right across a period (row), atomic radii decrease. This occurs because as more protons are added to the nucleus and more electrons are added within the same shell, there is no increased shielding between the protons and electrons, but there is increased attractive electrostatic force. This effect decreases the atomic radius. In contrast, as one moves from top to bottom down 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 forces and increasing the atomic radius. Because carbon and silicon are in the same group, and silicon is farther down the periodic table, it will have a larger atomic radius because of its extra electron shell. (C) and (D) are incorrect because all elements in the same group have the same number of valence electrons.
3. C
Atomic 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, III, is irrelevant, because it does not impact these attractive forces. Statements I and II are correct, and III is incorrect.
4. C
Ionization energy is related to the same set of forces that explains atomic radius, as well as the rules governing maintenance of a full valence shell octet. The first set of rules dictates that the stronger the attractive forces between the outer electron (the electron to be ionized) and the positively charged nucleus, the more energy will be required to ionize. As a result, strong attractive forces, which make the atomic radius smaller toward the right of a period or the top of a group, will also increase the first ionization energy. With this information alone, one could guess that ionization energies for beryllium (Be) should be higher than those for Li (lithium), eliminating (A) and (B). How should we choose between the two ionization energies for beryllium? The first ionization energy is usually dramatically lower than the second. This property holds true for the same reasons previously discussed. For example, upon removing one electron from beryllium, the ion is Be1+, which has one more proton in its nucleus than it has electrons surrounding it. Thus, there is a heightened electrostatic force between the positive nucleus and the now less-negative electron cloud, meaning that all remaining electrons will be held more tightly than before. Removing a second electron will be even more difficult than the initial electron removal, making the second ionization energy higher than the first. Furthermore, the first ionization energy for Li is 520.2 kJ/ mol, the first ionization energy for Be is 899.5 kJ/mol, and the second ionization energy for Be is 1,757.1 kJ/mol.
5. D
Selenium is on the right side of the periodic table (atomic number 34), too far right to be a metal or metalloid. However, it does not lie far enough to the right to fall under column 7A, which would classify it as a halogen. It is a nonmetal.
6. C
The trend in the periodic table demonstrated by the figure is correct for increasing electronegativity and first ionization energy. Electronegativity describes how strong an attraction an element will have for electrons in a bond. A nucleus with a stronger electrostatic pull due to its positive charge will have a higher electronegativity. An arrow pointing toward the right represents this because effective nuclear charge increases toward the right side of a period. This mirrors the trend for ionization energies, because 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 proximity of the outermost electrons to the positive inner nucleus increases, making the positive charge more effective at attracting new electrons in a chemical bond (which leads to higher electronegativity). Similarly, the more effective the positive nuclear charge, the higher the first ionization energy. Because I and III are both correct, we can choose answer choice (C).