Rucete ✏ AP Chemistry In a Nutshell
2. The Periodic Table - Practice Questions
This chapter introduces the structure and logic behind the periodic table, focusing on trends such as electronegativity, ionization energy, atomic and ionic radii, and the relationships between electron configurations and chemical behavior across periods and groups.
(Multiple Choice — Click to Reveal Answer)
1. Trends in the periodic table show that elements become more metallic in character from the top of a group to the bottom. Which of these is an element whose properties are opposite those of the element at the top of its group?
(A) Krypton
(B) Strontium
(C) Uranium
(D) Bismuth
Answer
(D) — Bismuth is a metal with properties opposite to nonmetals like nitrogen at the top of the group.
2. In which of the following atoms do the valence electrons feel the greatest effective nuclear charge?
(A) Ca
(B) K
(C) As
(D) Br
Answer
(D) — Br is furthest right in its period, leading to the highest effective nuclear charge.
3. In which choice below are the elements ranked in order of increasing first ionization energy?
(A) P, Cl, S, Al, Ar, Si
(B) Ar, Cl, S, P, Si, Al
(C) Al, Si, P, S, Cl, Ar
(D) Al, Si, S, P, Cl, Ar
Answer
(C) — Ionization energy generally increases across a period, with a slight exception for phosphorus and sulfur.
4. There are only two liquid elements at room temperature. One of these is:
(A) Krypton
(B) Bismuth
(C) Uranium
(D) Bromine
Answer
(D) — Bromine is a liquid at room temperature along with mercury.
5. The atom with the smallest radius is:
(A) Strontium
(B) Krypton
(C) Tin
(D) Bromine
Answer
(B) — Krypton has a high effective nuclear charge and small size due to full orbitals.
6. Metallic behavior is generally associated with:
(A) elements with low ionization energies
(B) elements with very negative electron affinities
(C) elements with small atomic radii
(D) elements with high electronegativities
Answer
(A) — Metals lose electrons easily, hence have low ionization energies.
7. One way to estimate the boiling point of Pd is to:
(A) average the boiling points of Rh and Ag
(B) average the boiling points of Ni and Pt
(C) average the boiling points of Ir and Cu
(D) average the boiling points of Co and Au
Answer
(B) — Ni and Pt surround Pd in the same group, making them good references.
8. In which of the following pairs is the element with the lower boiling point listed first?
(A) Na, Cs
(B) Te, Se
(C) P, N
(D) Ba, Sr
Answer
(D) — For metals, boiling point tends to decrease down a group.
9. In which of the following pairs is the first element expected to have a higher electronegativity than the second?
(A) O, P
(B) Cs, Rb
(C) I, Br
(D) Al, P
Answer
(A) — Oxygen is to the right and above phosphorus, leading to a higher electronegativity.
10. Which ion has the largest radius?
(A) Cl⁻
(B) K⁺
(C) S²⁻
(D) Ca²⁺
Answer
(C) — S²⁻ has gained 2 electrons and thus has the weakest pull per electron, giving it the largest radius.
11. The effective nuclear charge that an electron in the valence shell feels generally increases:
(A) from left to right across a period and down a group
(B) from left to right across a period and up a group
(C) from right to left across a period and down a group
(D) from left to right across a period and no change down a group
Answer
(D) — Effective nuclear charge increases across a period but does not change much down a group due to shielding.
12. Which of the following groups does not contain any metals?
(A) Xe, Hg, Ge, O
(B) Cl, Al, Si, Ar
(C) C, S, As, H
(D) Cu, P, Se, Kr
Answer
(C) — All are nonmetals or metalloids; none are metals.
13. What is the correct order of decreasing size of the following ions?
(A) P³⁻ > Cl⁻ > K⁺ > Ca²⁺
(B) K⁺ > Cl⁻ > P³⁻ > Ca²⁺
(C) Ca²⁺ > K⁺ > Cl⁻ > P³⁻
(D) K⁺ > Cl⁻ > Ca²⁺ > P³⁻
Answer
(A) — More negative charge = more repulsion = larger size. Cations are smaller.
14. Which pair of elements is expected to have the most similar properties?
(A) Potassium and lithium
(B) Sulfur and phosphorus
(C) Silicon and carbon
(D) Lithium and magnesium
Answer
(D) — Li and Mg are diagonally related and show similar chemical behavior.
15. Which of the following ions is isoelectronic with neon?
(A) Na⁺
(B) F⁻
(C) Mg²⁺
(D) All of the above
Answer
(D) — All have 10 electrons, like neon.
16. Which element has the most exothermic (most negative) electron affinity?
(A) Li
(B) F
(C) Be
(D) Na
Answer
(B) — Fluorine has the strongest tendency to gain electrons.
17. What causes atomic radius to decrease across a period?
(A) More neutrons
(B) Increased shielding
(C) Increased atomic mass
(D) Increased nuclear charge
Answer
(D) — More protons pull electrons closer, shrinking atomic size.
18. What type of elements are found along the "stair-step" line of the periodic table?
(A) Noble gases
(B) Transition metals
(C) Alkali metals
(D) Metalloids
Answer
(D) — Metalloids have properties of both metals and nonmetals and lie along the stair-step.
19. Which subatomic particle defines the identity of an element?
(A) Proton
(B) Neutron
(C) Electron
(D) Nucleus
Answer
(A) — The atomic number (number of protons) defines the element.
20. Which electron configuration represents a halogen?
(A) [Ne] 3s² 3p⁶
(B) [Ne] 3s² 3p⁵
(C) [He] 2s² 2p⁶
(D) [Ar] 4s¹
Answer
(B) — 3p⁵ corresponds to Group 17, the halogens.
21. Which particle is largest in size?
(A) F
(B) F⁻
(C) O²⁻
(D) Na⁺
Answer
(C) — More negative charge increases electron repulsion, making ions larger.
22. Which of the following is a correct trend across a period?
(A) Atomic radius increases
(B) Ionization energy decreases
(C) Electronegativity increases
(D) Metallic character increases
Answer
(C) — Across a period, atoms attract electrons more strongly → electronegativity increases.
23. What does a high binding energy in a PES spectrum indicate?
(A) A valence electron
(B) A weak nuclear attraction
(C) A strong nuclear attraction
(D) A large atomic radius
Answer
(C) — High binding energy means the electron is close to the nucleus and tightly held.
24. Which best explains why O has a smaller atomic radius than N?
(A) More core electrons
(B) Lower effective nuclear charge
(C) Higher effective nuclear charge
(D) Fewer protons
Answer
(C) — Oxygen has more protons than nitrogen, so it pulls its electrons in tighter.
25. What is the defining feature of noble gases?
(A) Full s and p sublevels
(B) One valence electron
(C) High metallic character
(D) Largest atomic radii in their periods
Answer
(A) — Noble gases have full outer electron shells, making them stable and unreactive.
26. Which process represents electron affinity correctly?
(A) X(g) → X⁺(g) + e⁻
(B) X⁺(g) → X⁺(aq)
(C) X⁺(g) + e⁻ → X(g)
(D) X(g) + e⁻ → X⁻(g)
Answer
(D) — Electron affinity refers to the energy change when a neutral atom gains an electron.
27. Which of the following elements has the most exothermic electron affinity?
(A) Li
(B) F
(C) Be
(D) Na
Answer
(B) — Fluorine is the most electronegative element and has the highest (most negative) electron affinity.
28. Which best explains the large jump in ionization energy after removal of the valence electrons?
(A) The second valence electron is harder to remove
(B) The nucleus gets stronger
(C) The next electron is from a core level
(D) Electrons move closer to the nucleus
Answer
(C) — Core electrons are more tightly bound; removing them requires much more energy.
29. What causes ionization energy to decrease down a group?
(A) Lower atomic mass
(B) More protons
(C) Increased shielding and distance from nucleus
(D) Higher electronegativity
Answer
(C) — Outer electrons are farther from the nucleus and more shielded, making them easier to remove.
30. Which of the following statements is true regarding atomic radii?
(A) Atomic radius increases left to right in a period
(B) Atomic radius decreases down a group
(C) Atomic radius increases down a group and decreases across a period
(D) Atomic radius is the same for all isotopes
Answer
(C) — Atomic radius increases down a group due to added shells and decreases across a period due to increasing nuclear charge.
31. What is the electron configuration for the most common ion of aluminum?
(A) 1s² 2s² 2p⁶ 3s² 3p¹
(B) 1s² 2s² 2p⁶
(C) 1s² 2s² 2p⁶ 3s²
(D) 1s² 2s² 2p⁶ 3s² 3p⁵
Answer
(B) — Al³⁺ loses all three of its valence electrons, leaving a configuration like neon.
32. Which pair of elements has the most similar chemical behavior?
(A) Li and Mg
(B) S and Cl
(C) Mg and Ca
(D) Si and N
Answer
(C) — Mg and Ca are both alkaline earth metals in the same group, showing similar reactivity.
33. Which of the following elements is a metalloid?
(A) Arsenic
(B) Aluminum
(C) Calcium
(D) Sulfur
Answer
(A) — Arsenic is located along the stair-step line and exhibits properties of both metals and nonmetals.
34. Which has the greatest effective nuclear charge?
(A) Na
(B) Mg
(C) Al
(D) Cl
Answer
(D) — Cl is furthest right in its period, with the most protons for its energy level and little added shielding.
35. Which best explains why Mg has a higher first ionization energy than Na?
(A) Mg has a greater number of protons
(B) Mg is more massive
(C) Mg has more neutrons
(D) Mg is a noble gas
Answer
(A) — More protons in Mg increase the attraction to the electrons, making them harder to remove.
36. Explain the periodic trend in atomic radius as you move from left to right across a period.
Answer
Atomic radius decreases from left to right due to increasing effective nuclear charge, which pulls electrons closer to the nucleus.
37. Why do elements in the same group have similar chemical properties?
Answer
Elements in the same group have the same number of valence electrons, leading to similar reactivity and chemical properties.
38. What is meant by “isoelectronic,” and give an example from the periodic table.
Answer
Isoelectronic species have the same electron configuration. For example, Na⁺, Mg²⁺, and Ne all have 10 electrons.
39. Describe the trend in ionization energy as you go down a group and explain why this trend occurs.
Answer
Ionization energy decreases down a group because electrons are farther from the nucleus and more shielded, making them easier to remove.
40. State the relationship between effective nuclear charge and atomic radius.
Answer
As effective nuclear charge increases, atomic radius decreases because electrons are pulled closer to the nucleus.
41. Why does F have a higher electron affinity than O?
Answer
F has more protons, so its nucleus attracts extra electrons more strongly, resulting in a more negative electron affinity.
42. Explain why removing an electron from a core (inner) shell requires much more energy than removing one from a valence shell.
Answer
Core electrons are much closer to the nucleus and experience less shielding, so they are held more tightly and require more energy to remove.
43. What defines a metalloid and name two examples.
Answer
Metalloids have properties intermediate between metals and nonmetals; examples: silicon (Si), arsenic (As).
44. What is the weighted average atomic mass and how is it calculated?
Answer
It’s the average mass of all naturally occurring isotopes, calculated by multiplying each isotope’s mass by its abundance and summing the results.
45. Why are anions larger than their neutral atoms?
Answer
Adding electrons increases electron-electron repulsion and decreases the effective nuclear charge per electron, expanding the size of the ion.
46. Why do noble gases have very low reactivity?
Answer
Noble gases have full valence shells, making them stable and unlikely to gain or lose electrons.
47. How can photoelectron spectroscopy (PES) be used to determine electron configuration?
Answer
PES measures the energies of electrons ejected from different orbitals, allowing the number of electrons in each orbital to be inferred from peak intensities.
48. Explain the difference between atomic number and mass number.
Answer
Atomic number is the number of protons; mass number is the total number of protons and neutrons in the nucleus.
49. What is a diagonal relationship in the periodic table? Give an example.
Answer
A diagonal relationship refers to similarities in properties between elements diagonally adjacent in the periodic table, such as Li and Mg.
50. Give an example and brief explanation of why a mass number listed in the periodic table is often not a whole number.
Answer
Atomic masses are weighted averages of all isotopes. For example, chlorine’s atomic mass is about 35.5 because it is a mix of Cl-35 and Cl-37.