Study guide
7.1 Atomic & Nuclear Structure, Isotopes
NAPT Navy Advanced Programs Test 2026
Track your progress as you read
Finished sections get a check mark.
1 Introduction
About the NAPT & Navy Nuclear Field Program NFa & NFb Qualification Routes and Scoring How to Use This Guide
2 Chapter 1: Mathematics I — Arithmetic & Algebra
1.1 Arithmetic, Fractions, Decimals & Percentages 1.2 Ratios, Proportions & Unit Conversions 1.3 Algebra I: Linear Equations & Inequalities 1.4 Algebra II: Exponents, Radicals & Quadratic Equations
3 Chapter 2: Mathematics II — Geometry, Trigonometry & Word Problems
2.1 Geometry: Area, Volume & Right Triangles 2.2 Trigonometry Basics 2.3 Word Problems & Multi-Step Reasoning
4 Chapter 3: Physics I — Mechanics & Energy
3.1 Kinematics & Motion 3.2 Forces & Newton's Laws 3.3 Work, Energy & Power
5 Chapter 4: Physics II — Thermodynamics, Electricity & Waves
4.1 Thermodynamics & Heat Transfer 4.2 Electricity: Ohm's Law & Series/Parallel Circuits 4.3 Kirchhoff's Laws & Magnetism 4.4 Waves, Fluids & Pressure
6 Chapter 5: Chemistry I — Atomic Structure & Reactions
5.1 Atomic Structure & the Periodic Table 5.2 Chemical Bonding & Reactions 5.3 Stoichiometry & the Mole Concept
7 Chapter 6: Chemistry II — Gas Laws, Solutions & Thermochemistry
6.1 Gas Laws 6.2 Solutions, Acids & Bases 6.3 States of Matter & Thermochemistry
8 Chapter 7: Nuclear Science
7.1 Atomic & Nuclear Structure, Isotopes 7.2 Radioactivity, Decay & Half-Life 7.3 Nuclear Fission, Chain Reactions & Critical Mass 7.4 Reactor Basics & Radiation Safety
9 Chapter 8: Administration Boundaries & Editorial Review
8.1 Administration Boundaries & Personal Readiness 8.2 Integrated Review & Local Practice
7.1 Atomic & Nuclear Structure, Isotopes
Key Takeaways Explain This Topic AI
Atomic number (Z) equals the number of protons and defines the element; isotopes of that element share Z but differ in neutron number (N) and mass number (A)
Isotope notation follows N = A − Z; for example, Uranium-235 (Z = 92, A = 235) has 143 neutrons, while Uranium-238 has 146
Nuclear binding energy comes from the mass defect — the mass that seems to disappear when separate nucleons bind into a nucleus — converted to energy via E = mc²
Binding energy per nucleon reaches a broad maximum in the iron–nickel region; Nickel-62 is commonly listed as the highest, with Iron-56 very close.
Because c² is such a large number, a minuscule nuclear mass defect converts into far more energy than any chemical reaction releases from the same mass of material
Last updated: August 2026
Atomic and nuclear structure is included as public, textbook-level STEM enrichment. The Navy's public NAPT qualification rules do not publish a nuclear-science domain or content blueprint, so this chapter must not be read as required exam scope.
Nucleons: The Building Blocks of the Nucleus
Every atom has a tiny, dense core called the nucleus , made up of two types of particles collectively called nucleons :
- Particle Charge Relative Mass Location
- Proton +1 ~1 atomic mass unit (amu) Inside the nucleus
- Neutron 0 (neutral) ~1 amu (slightly more than a proton) Inside the nucleus
- Electron −1 ~1/1,836 amu (negligible) Orbiting the nucleus
Protons and neutrons together account for essentially all of an atom's mass, since electrons are thousands of times lighter. The number of protons is what makes an atom a particular chemical element — change the proton count and you change the element entirely.
Atomic Number, Mass Number, and Neutron Number
Three numbers describe any nucleus:
Atomic number (Z) — the number of protons. This defines the element (every carbon atom has Z = 6; every uranium atom has Z = 92).
Mass number (A) — the total number of protons plus neutrons (nucleons).
Neutron number (N) — simply A minus Z.
This gives a core relationship used repeatedly in the local examples:
- N = A − Z
- Isotopes and Isotope Notation
Isotopes are atoms of the same element (same Z, same number of protons) that have different numbers of neutrons (different N and therefore different A). Isotopes are written in one of two equivalent ways: the full notation with the mass number as a left superscript and the atomic number as a left subscript in front of the element symbol, or — far more common in test questions and everyday nuclear terminology — as Element-A, such as Carbon-14 or Uranium-235.
- Isotope Common Name Protons (Z) Neutrons (N) Mass Number (A)
- Hydrogen-1 Protium 1 0 1
- Hydrogen-2 Deuterium 1 1 2
- Hydrogen-3 Tritium 1 2 3
- Carbon-12 Stable carbon 6 6 12
- Carbon-14 Radioactive carbon 6 8 14
- Uranium-235 Fissile uranium 92 143 235
- Uranium-238 Fertile, not readily fissile 92 146 238
Notice that hydrogen's three isotopes all have exactly one proton (that is what makes them hydrogen) but zero, one, and two neutrons respectively. The same logic applies to the two uranium isotopes that matter most for nuclear science: U-235 and U-238 both have 92 protons, but U-235 has three fewer neutrons — a difference that turns out to matter enormously in section 7.3, because U-235 is the isotope that sustains a chain reaction easily and U-238 does not.
Nuclear Binding Energy (Conceptual)
Here is a question worth sitting with: protons all carry the same positive charge, so they should repel each other strongly at the tiny distances inside a nucleus. What holds the nucleus together?
The answer is the strong nuclear force — a fundamental force that is far stronger than electromagnetic repulsion, but only over extremely short distances (roughly the width of a nucleus). At nuclear distances, the strong force overpowers the electrostatic repulsion between protons and binds protons and neutrons together into a stable nucleus.
Holding the nucleus together this way has a measurable consequence: if you added up the mass of every proton and neutron separately, the total would be slightly more than the mass of the actual bound nucleus. That difference is called the mass defect (Δm) — mass that seems to disappear when the nucleons bind together. It has not disappeared; it has been converted into the energy that holds the nucleus together, called the nuclear binding energy .
Mass-Energy Equivalence (E = mc²) — Conceptually
The mass defect converts into binding energy according to Albert Einstein's mass-energy equivalence relationship:
E = mc²
This local overview emphasizes the concept rather than numerical calculation: E (energy) and m (mass) are two forms of the same underlying quantity, and they convert into each other through the constant c² (the speed of light, squared — an enormous number, about 9 × 10¹⁶ m²/s²). Because c² is so large, converting even a tiny amount of mass releases a tremendous amount of energy.
This is exactly why nuclear reactions release far more energy, per gram of material, than chemical reactions like burning fuel. A chemical reaction only rearranges the outer electrons of atoms — the nuclei, and their mass, stay untouched. A nuclear reaction changes the nucleus itself, converting a measurable fraction of mass directly into energy. That difference in scale is the entire reason nuclear power and nuclear weapons exist as distinct categories from chemical power and chemical explosives.
Binding Energy Per Nucleon: The Iron–Nickel Region
If you divide total binding energy by the number of nucleons in a nucleus, you get the binding energy per nucleon —a rough measure of how tightly bound a nucleus is. Plotted against mass number, the curve reaches a broad maximum in the iron–nickel region . Nickel-62 is commonly listed as having the highest binding energy per nucleon, while Iron-56 is extremely close; simplified discussions often call this the iron peak.
That single curve previews the rest of this chapter:
Fusion of sufficiently light nuclei can release energy when products are more tightly bound and move toward the iron–nickel region.
Fission of very heavy nuclei such as uranium can release energy when fragments are more tightly bound and move toward that region; section 7.3 gives a conceptual overview.
Applying the Notation
The following local textbook examples ask you to extract Z, N, or A from an isotope's name or notation:
Example 1: How many neutrons does Uranium-235 have? Z = 92 (uranium's atomic number never changes). A = 235. N = A − Z = 235 − 92 = 143 neutrons .
Example 2: Carbon-14 is used in radiocarbon dating. How many neutrons does it have? Z = 6 (carbon). A = 14. N = 14 − 6 = 8 neutrons (two more than ordinary Carbon-12, which has 6).