Quiz: Periodic Table and Periodic Trends — 28 questions

1. Which principle determines the order of elements in the modern periodic table?

A. Increasing atomic mass

B. Decreasing atomic number

C. Decreasing atomic mass

D. Increasing atomic number

Answer: (D)

Explanation

The modern table orders all 118 elements by increasing atomic number, which reflects the number of protons. Atomic mass was used in earlier arrangements, including Mendeleev’s table, but is not the modern basis.

2. What significant contribution did Dmitri Mendeleev make in 1869?

A. He arranged 63 elements by atomic mass and left gaps for undiscovered elements.

B. He observed that every eighth element repeated the properties of the first.

C. He measured atomic numbers with X-ray emission and established the modern Periodic Law.

D. He grouped elements into triads whose middle atomic weight was approximately averaged.

Answer: (A)

Explanation

Mendeleev organized 63 elements by increasing atomic mass, deliberately leaving gaps and predicting the properties of missing elements. X-ray measurements belong to Moseley, while triads and the eighth-element pattern belong to earlier researchers.

3. How did Moseley’s 1913 research lead to the modern Periodic Law?

A. By using atomic weights to group elements into sets of three related elements

B. By arranging known elements by mass and predicting properties of missing elements

C. By using X-ray emission to determine atomic numbers and reorder elements accordingly

D. By observing repeated properties after every eighth element in a mass-based sequence

Answer: (C)

Explanation

Moseley used X-ray emission to determine exact atomic numbers and showed that elements should be arranged by those numbers. The other approaches describe Döbereiner, Mendeleev, and Newlands, respectively.

4. What do the periods and groups of the modern periodic table represent?

A. Seven horizontal rows and eighteen vertical columns

B. Eighteen horizontal rows and seven diagonal columns

C. Eighteen horizontal rows and seven vertical columns

D. Seven diagonal rows and eighteen vertical columns

Answer: (A)

Explanation

The table contains seven horizontal periods and eighteen vertical groups. Confusing the orientations or the counts reverses the table’s basic structure.

5. Why do elements in the same group generally have similar chemical properties?

A. They have identical physical properties throughout the group.

B. They undergo the same gradual change from left to right.

C. They have the same number of valence electrons.

D. They occupy the same horizontal period.

Answer: (C)

Explanation

Shared valence-electron counts give elements within a group similar chemical behavior, although physical properties vary gradually down the group. Horizontal periods instead describe left-to-right trends, not shared group chemistry.

6. An element tends to lose electrons and form a positive ion, while another tends to gain electrons and form a negative ion; how should they be classified?

A. Both are metalloids because they form ions.

B. The first is a metal, and the second is a nonmetal.

C. The first is a nonmetal, and the second is a metal.

D. The first is a metalloid, and the second is a metal.

Answer: (B)

Explanation

Metals commonly lose electrons to form positive ions, whereas nonmetals commonly gain electrons to form negative ions. Metalloids combine characteristics of metals and nonmetals rather than fitting both ion-forming patterns directly.

7. Which description correctly matches the four blocks of the periodic table?

A. The s-block contains transition elements, the d-block involves s subshells, the p-block contains lanthanides, and the f-block covers Groups 13–18.

B. The s-block contains lanthanides and actinides, the d-block covers Groups 13–18, the p-block involves s subshells, and the f-block contains transition elements.

C. The s-block covers Groups 13–18, the d-block contains lanthanides and actinides, the p-block contains transition elements, and the f-block involves s subshells.

D. The s-block involves s subshells, the d-block contains transition elements, the p-block covers Groups 13–18, and the f-block contains lanthanides and actinides.

Answer: (D)

Explanation

Block names reflect the subshell occupied by valence electrons, with the d-block corresponding to transition elements and the f-block to the two bottom series. The alternatives interchange these block assignments.

8. Which feature distinguishes alkali metals from alkaline earth metals?

A. Alkali metals are Group 18 elements with complete outer shells and low reactivity.

B. Alkali metals are Group 2 elements with two valence electrons and lower reactivity.

C. Alkali metals are Group 1 elements with one valence electron and greater reactivity.

D. Alkali metals are Group 17 elements that accept one electron and form salts.

Answer: (C)

Explanation

Group 1 alkali metals have one valence electron, react strongly with water, and are more reactive than Group 2 alkaline earth metals. The other descriptions refer to alkaline earth metals, halogens, and noble gases.

9. Which pairing correctly identifies a highly reactive nonmetal group and its typical behavior?

A. Halogens in Group 17 readily accept one electron and form salts.

B. Alkali metals in Group 1 readily accept one electron and form salts.

C. Noble gases in Group 18 readily accept one electron and form salts.

D. Alkaline earth metals in Group 2 readily accept one electron and form salts.

Answer: (A)

Explanation

Halogens are highly reactive Group 17 nonmetals that readily gain one electron and form salts. Noble gases have complete outer shells, while Groups 1 and 2 are metals that tend to lose electrons.

10. Why are noble gases generally almost unreactive under normal conditions?

A. They readily accept one electron to form salts.

B. They have one valence electron available for easy loss.

C. Their outer electron shells are complete.

D. They contain two valence electrons and react more slowly than alkali metals.

Answer: (C)

Explanation

Complete outer shells make noble gases chemically stable and therefore nearly unreactive under normal conditions. Their stability does not result from having one or two readily transferable valence electrons, and salt formation is characteristic of halogens.

11. What does the period number indicate in the periodic table for a main-group element?

A. The number of valence electrons

B. The number of neutrons in the nucleus

C. The number of electrons in the outer orbital

D. The number of occupied electron shells

Answer: (D)

Explanation

The period identifies how many electron shells are occupied. The group number, not the period number, indicates the valence-electron count for relevant main-group elements.

12. Which electron configuration belongs to an element in Group 13 and Period 3?

A. 1s² 2s² 2p⁶ 3s² 3p³

B. 1s² 2s² 2p⁶ 3s² 3p¹

C. 1s² 2s² 2p⁶ 3s¹

D. 1s² 2s² 2p⁶ 3s²

Answer: (B)

Explanation

A Period 3, Group 13 element has three occupied shells and three valence electrons, giving the stated 3s²3p¹ ending. The other configurations represent different valence-electron patterns.

13. How is atomic radius commonly defined for a neutral element?

A. Half the diameter of an isolated atomic nucleus

B. Half the distance between two identical bonded atoms

C. The distance from the nucleus to the outermost electron

D. The distance between two oppositely charged ions

Answer: (B)

Explanation

Atomic radius is conventionally measured as half the internuclear distance between two identical bonded atoms. The other descriptions refer to informal atomic size ideas or different distances.

14. Why does atomic radius generally decrease from left to right across a period?

A. Greater nuclear charge pulls the electron cloud closer

B. Valence electrons occupy progressively higher principal shells

C. The nucleus loses protons as atomic number increases

D. Additional shells shield the outer electrons more strongly

Answer: (A)

Explanation

Across a period, increasing nuclear charge attracts the electron cloud more strongly while electrons are added within the same general shell. Additional-shell shielding is the trend associated with moving down a group.

15. What best explains why atomic radius generally increases down a group?

A. Additional electron shells increase shielding and expand the atom

B. Electron pairing removes repulsion from the outer shell

C. Valence electrons move into lower principal shells

D. Fewer protons reduce the nuclear attraction across the group

Answer: (A)

Explanation

Moving down a group adds occupied shells, increasing shielding and placing valence electrons farther from the nucleus. The nucleus generally gains protons, so reduced proton number is not the explanation.

16. Compared with its neutral atom, what size change generally occurs when an atom forms a cation?

A. Its radius increases because the positive charge adds an electron shell

B. Its radius increases because fewer electrons reduce nuclear attraction

C. Its radius remains unchanged because the nucleus is unaffected

D. Its radius decreases because electron loss strengthens attraction on remaining electrons

Answer: (D)

Explanation

A cation is generally smaller because losing electrons can remove a shell and increases the nucleus’s attraction on the remaining electrons. A larger radius is generally associated with anions formed by electron gain.

17. How does first ionization energy generally vary across the periodic table?

A. It decreases down a group and increases from left to right across a period

B. It decreases in both directions because atoms become more reactive

C. It increases down a group and decreases from left to right across a period

D. It increases in both directions because nuclear charge always dominates

Answer: (A)

Explanation

Larger size and shielding make electrons easier to remove down a group, while increasing effective nuclear charge makes removal harder across a period. The opposite trend reverses both established patterns.

18. Which order correctly ranks the first ionization energies of Group 1 elements from highest to lowest?

A. Li > K > Na > Cs > Rb

B. Na > Li > Rb > K > Cs

C. Cs > Rb > K > Na > Li

D. Li > Na > K > Rb > Cs

Answer: (D)

Explanation

First ionization energy decreases down Group 1 as atomic size and shielding increase, producing the stated sequence. Cesium therefore has a lower value than lithium, not a higher one.

19. What does the first electron affinity of an element measure?

A. The enthalpy change when gaseous negative ions gain electrons to form gaseous dinegative ions

B. The enthalpy change when gaseous atoms gain electrons to form gaseous uninegative ions

C. The energy required to remove electrons from gaseous atoms to form positive ions

D. The attraction between bonded atoms and a shared pair of electrons in a molecule

Answer: (B)

Explanation

First electron affinity concerns adding one electron to each neutral gaseous atom and forming gaseous uninegative ions. Adding an electron to a negative ion describes second electron affinity instead.

20. Why is oxygen’s second electron affinity positive while its first electron affinity is negative?

A. The second electron enters a lower-energy shell closer to the nucleus

B. The oxygen atom loses an electron before the second electron is added

C. The incoming electron is repelled by the already negative oxygen ion

D. The neutral oxygen atom repels the first incoming electron more strongly

Answer: (C)

Explanation

The second electron must be added to a negative ion, so electron–electron repulsion requires energy input. The first addition to a neutral oxygen atom releases energy, giving a negative enthalpy change.

21. How does electron affinity generally change across a period and down a group?

A. It becomes less negative across a period and increases down a group

B. It becomes more negative across a period and increases down a group

C. It becomes less negative across a period and decreases down a group

D. It becomes more negative across a period and decreases down a group

Answer: (D)

Explanation

Across a period, electron affinity generally becomes more negative, whereas it decreases down a group as atomic size and nuclear distance increase. The opposing trend in the second option reverses both relationships.

22. Which statement correctly describes electronegativity on the Pauling scale?

A. It measures ionization energy, with fluorine rated 0.8

B. It measures an atom’s ability to attract shared electrons, with fluorine rated 4.0

C. It measures an atom’s tendency to gain electrons, with fluorine rated 0.8

D. It measures bond strength, with alkali metals rated 4.0

Answer: (B)

Explanation

Electronegativity describes attraction for a shared pair of bonding electrons, and fluorine has the maximum Pauling value of 4.0. The alternatives confuse it with electron affinity, bond strength, or ionization energy.

23. What products form when sodium reacts with water?

A. Solid sodium oxide and liquid hydrogen

B. Aqueous sodium chloride and oxygen gas

C. Aqueous sodium hydroxide and hydrogen gas

D. Solid sodium peroxide and aqueous hydrogen chloride

Answer: (C)

Explanation

Sodium reacts vigorously with water to produce sodium hydroxide solution and hydrogen gas. Sodium oxide, sodium peroxide, and sodium chloride arise in different reactions.

24. Why are Period Three oxides and chlorides generally more covalent from Group 4 onward?

A. Atomic size increases across the period, making ionic bonding more favorable

B. Electronegativity increases across the period, strengthening covalent character

C. Metallic character increases across the period, producing more mobile ions

D. Chlorine and oxygen become less electronegative across the period, favoring electron transfer

Answer: (B)

Explanation

Increasing electronegativity across Period Three makes bonding in Groups 4–7 compounds more covalent. The other choices reverse periodic trends or incorrectly imply greater ionic character.

25. Which description identifies a basic oxide?

A. It reacts with both acids and bases, as aluminium oxide does

B. It hydrolyzes in water to produce an acidic solution, as many Period Three chlorides do

C. It reacts with water to form an acid and is generally a covalent nonmetal oxide

D. It reacts with water to form an alkali and is usually an ionic metal oxide

Answer: (D)

Explanation

Basic oxides, such as sodium oxide and calcium oxide, commonly react with water to form alkaline solutions and are usually ionic. The second description defines acidic oxides, while the others describe amphoteric oxides or acidic chloride hydrolysis.

26. Which substance is an example of an acidic oxide?

A. Calcium oxide, which is a generally ionic metal oxide

B. Aluminium oxide, which reacts with both acids and bases

C. Sodium oxide, which can react with water to produce an alkali

D. Sulfur dioxide, which can react with water to produce an acid

Answer: (D)

Explanation

Sulfur dioxide is a covalent nonmetal oxide that forms an acid when it reacts with water. Sodium oxide and calcium oxide are basic, while aluminium oxide is amphoteric.

27. A laboratory sample reacts with both hydrochloric acid and sodium hydroxide. Which Period Three oxide best matches this behavior?

A. Sodium oxide

B. Sulfur trioxide

C. Aluminium oxide

D. Phosphorus pentoxide

Answer: (C)

Explanation

Aluminium oxide is amphoteric because it reacts with both acids and bases. Sodium oxide is basic, whereas sulfur trioxide and phosphorus pentoxide are acidic oxides.

28. Which statement accurately compares Period Three chlorides in water?

A. All Period Three chlorides form alkaline solutions because metal chlorides release hydroxide ions

B. Group 1 and Group 2 chlorides are generally neutral, while aluminium-to-sulfur chlorides form acidic solutions

C. Group 1 and Group 2 chlorides form acidic solutions, while aluminium-to-sulfur chlorides remain neutral

D. All Period Three chlorides remain neutral because chloride ions do not hydrolyze

Answer: (B)

Explanation

Group 1 and Group 2 chlorides are generally neutral in water, whereas chlorides from aluminium to sulfur hydrolyze and produce acidic solutions. The alternatives reverse or deny this distinction and incorrectly predict alkaline behavior.