5.1 Atomic Structure & the Periodic Table

Every chemistry question on the Navy Advanced Programs Test (NAPT) builds on the same foundation: what an atom is made of, and how the periodic table organizes the elements that make up everything in the physical world — including the fuel, coolant, alloys, and reactor materials sailors work with every day. This section covers atomic structure, isotopes, the periodic table's organization, the first 20 elements, common oxidation states, and basic electron configuration — the vocabulary and number sense that every later chemistry section (bonding, reactions, gas laws, solutions) depends on.

Protons, Neutrons & Electrons

Every atom is built from three subatomic particles:

Protons and neutrons together make up the dense, positively charged nucleus at an atom's center; electrons occupy the much larger surrounding space in shells (energy levels). Because protons and electrons carry equal and opposite charge, a neutral atom always has the same number of protons as electrons.

Atomic Number vs. Mass Number

Two numbers describe every atom:

• Atomic number (Z) — the number of protons in the nucleus. This number defines which element the atom is; every atom of a given element has the same atomic number, and no two elements share one.
• Mass number (A) — the total number of protons plus neutrons in the nucleus (electrons are too light to matter). Mass number is what you'd round an atom's mass to in whole atomic mass units.

For example, a neutral atom of potassium always has atomic number 19 (19 protons). If that same atom has 20 neutrons, its mass number is 19 + 20 = 39, written as potassium-39.

Isotopes

Isotopes are atoms of the same element (same atomic number, same proton count) that have different numbers of neutrons — and therefore different mass numbers. Carbon, for instance, has three naturally occurring isotopes:

All three are chemically carbon — same 6 protons, same chemical behavior — but they have different masses and, in the case of carbon-14, different nuclear stability. (Radioactive decay of unstable isotopes is covered in depth in the Nuclear Science chapter.)

When an atom gains or loses electrons rather than neutrons, it becomes an ion — a charged particle — rather than an isotope. Sodium (Na) losing its outer electron becomes Na⁺ (11 protons, 10 electrons); chlorine (Cl) gaining an electron becomes Cl⁻ (17 protons, 18 electrons). Ions are the building blocks of ionic bonding, covered in the next section.

Periodic Table Organization

The periodic table arranges all known elements by increasing atomic number into 7 periods (horizontal rows) and 18 groups (vertical columns).

• Periods correspond to the number of electron shells an atom has. Elements in Period 1 (H, He) use only one shell; elements in Period 3 (Na through Ar) use three.
• Groups collect elements with the same number of valence (outermost-shell) electrons, which gives them similar chemical behavior. This is why, for example, all Group 1 elements react vigorously with water and all Group 17 elements form salts with metals.

Key groups to know by name:

The table is also divided by broad category: metals (most of the left and center of the table) tend to lose electrons and conduct electricity; nonmetals (upper right) tend to gain electrons; and metalloids — boron, silicon, germanium, arsenic, antimony, and tellurium — sit along the "staircase" between them and share properties of both.

The First 20 Elements

The NAPT's chemistry content leans heavily on the first 20 elements — you should be able to recall their symbols and atomic numbers cold:

Notice the pattern: elements 3, 11, and 19 (Li, Na, K) are all alkali metals; elements 9 and 17 (F, Cl) are both halogens; elements 2, 10, and 18 (He, Ne, Ar) are all noble gases. Recognizing these families by atomic number — not just by memorized symbol — is what actually gets tested.

Common Oxidation States

An oxidation state (or oxidation number) describes the charge an atom effectively has in a compound, based on how many electrons it has gained, lost, or unequally shared. For the first 20 elements, oxidation state tracks closely with group number:

Hydrogen is a special case: it usually shows +1 (as in H2O or HCl) but takes -1 in metal hydrides such as NaH, where it bonds to an even less electronegative metal.

Electron Configuration Basics

Electrons fill an atom's shells from the lowest energy level outward, a pattern called the Aufbau principle ("building-up" principle). Two equivalent ways to describe this show up on the NAPT:

1. Simplified shell (Bohr) model — count electrons per shell, capped at 2 in the first shell and 8 in the second and third shells for the elements covered here:

2. Standard subshell notation — write out sublevels (s, p) in order of filling:

• Hydrogen (Z=1): 1s¹
• Oxygen (Z=8): 1s² 2s² 2p⁴
• Neon (Z=10): 1s² 2s² 2p⁶
• Sodium (Z=11): 1s² 2s² 2p⁶ 3s¹ (shorthand: [Ne] 3s¹)
• Chlorine (Z=17): 1s² 2s² 2p⁶ 3s² 3p⁵ (shorthand: [Ne] 3s² 3p⁵)

The outermost shell's electrons — the valence electrons — are what actually participate in bonding. Atoms are most chemically stable with a full outer shell of 8 valence electrons (2 for hydrogen and helium), a pattern called the octet rule. Sodium's single valence electron is loosely held and easily lost; chlorine's 7 valence electrons need just one more to complete an octet, so it's easily gained. That imbalance — one atom wanting to give an electron away, another wanting to take one — is exactly what drives the ionic and covalent bonding covered in the next section.

Key Takeaways

• Atomic number (Z) = protons; mass number (A) = protons + neutrons; isotopes share Z but differ in neutron count and A.
• The periodic table's 18 groups share valence-electron counts and chemical behavior; its 7 periods track electron shell count.
• Know the first 20 elements (H through Ca) by atomic number, and recognize the alkali metals, halogens, and noble gases among them.
• Oxidation state tracks group number: +1 (Group 1), +2 (Group 2), -1 (Group 17, halogens), 0 (Group 18, noble gases).
• Valence electrons drive reactivity; atoms bond to satisfy the octet rule of 8 stable outer electrons.