Question 1
Sodium (Na) is in Group 1 and has electron configuration . A small piece of sodium reacts vigorously with water and forms in the process. How does sodium’s atomic structure explain this chemical behavior?
- Sodium has a full valence shell, so it does not need to react and stays uncharged.
- Sodium has 1 valence electron, which it loses easily to reach a stable noble-gas configuration, making it very reactive and forming .
- Sodium’s high reactivity is mainly because it has 11 total electrons, and elements with more electrons always react more.
- Sodium forms because it gains one electron to fill its 3rd shell.
Explanation: This question tests your understanding of how atomic structure—particularly the number of valence electrons—determines chemical behavior including reactivity, bonding tendency, and ion formation. The number of valence electrons (electrons in the outermost shell) is THE key structural feature that determines how an element behaves chemically: atoms with 1-3 valence electrons (groups 1, 2, 13—metals) tend to LOSE those electrons easily because achieving a full inner shell (matching the previous noble gas) is energetically favorable, making these elements reactive metals that form positive ions. Atoms with 5-7 valence electrons (groups 15, 16, 17—nonmetals) tend to GAIN electrons to complete their outer shells to 8 (matching the next noble gas), making these reactive nonmetals that form negative ions. Atoms with 8 valence electrons (noble gases) are already stable and don't react under normal conditions because they already have full outer shells—nothing to gain by reacting! This is why elements in the same group (same valence electron count) show similar chemical behavior: all group 1 elements are reactive metals forming +1 ions, all group 17 elements are reactive nonmetals forming -1 ions. For sodium, with its 1 valence electron in the 3s orbital, the structure promotes high reactivity because losing that single electron exposes a full inner shell like neon, allowing it to form Na⁺ during vigorous reactions like with water. Choice B correctly relates atomic structure (valence electrons, configuration, or periodic position) to chemical behavior using sound cause-effect reasoning by explaining how sodium's 1 valence electron drives it to lose that electron for stability, leading to reactivity and Na⁺ formation. Choice A fails because sodium does not have a full valence shell—its 3s¹ is incomplete, so it must react to achieve stability, not stay uncharged; supportive correction: elements with full shells like noble gases are unreactive, but sodium's structure demands electron loss. The structure-to-behavior prediction framework: (1) Determine valence electrons from group number or configuration: Group 1 = 1 valence, Group 2 = 2 valence, Group 13 = 3 valence, Group 14 = 4 valence, Group 15 = 5 valence, Group 16 = 6 valence, Group 17 = 7 valence, Group 18 = 8 valence (or 2 for helium). (2) Apply the valence rules: 1-3 valence = metal behavior (lose electrons, form positive ions, reactive if few valence), 5-7 valence = nonmetal behavior (gain electrons, form negative ions, reactive if near 8), 8 valence = noble gas behavior (stable, unreactive, no ions). (3) Predict specifics: Number of valence often equals bonds formed (carbon's 4 valence → forms 4 bonds usually). Group number predicts ion charge (group 1 → +1 ion from losing 1 valence electron). Reactivity extremes at group 1 (most reactive metals) and group 17 (most reactive nonmetals). Valence electron thinking: imagine you're an atom with 1 valence electron (like sodium). You could either (a) gain 7 more electrons to fill your shell to 8 (hard! requires 7 new electrons), or (b) lose that 1 electron to reveal the full shell underneath (easy! just remove 1). Option (b) wins—lose the 1 electron, form Na⁺, match neon's stability. Now imagine you're an atom with 7 valence electrons (like chlorine). You could (a) lose all 7 to reveal inner shell (hard! removing 7 electrons), or (b) gain 1 more to complete your octet to 8 (easy! just add 1). Option (b) wins—gain 1 electron, form Cl⁻, match argon. This thought experiment explains why metals lose electrons and nonmetals gain them: whichever path requires fewer electron changes wins!