Chemistry 2e · Atoms, Molecules, and Ions
Atomic Structure and Symbolism
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In 30 seconds
Now that the atom is known to contain protons, neutrons, and electrons, chemists need a compact way to say exactly which atom they mean. Two numbers do the job. The atomic number (Z) is the number of protons — it defines the element: 6 protons means carbon, 92 means uranium, no exceptions. The mass number (A) is the total number of protons plus neutrons. Because atoms of the same element can have different numbers of neutrons, the same element exists as several isotopes, each written with a Nuclide A specific atom with defined Z and A Full entry → symbol like 126C or 23592U. This topic also covers ions (atoms with unequal proton and electron counts) and the Average atomic mass Abundance-weighted mean of isotope masses Full entry → — the weighted average of Isotope Atom of an element with a different neutron count Full entry → masses that appears as the decimal number on the periodic table. Most of what you will do in later chapters — stoichiometry, nuclear chemistry, mass spectrometry — depends on these symbols and numbers.
Why this matters
The atomic number and mass number are the keys that unlock the periodic table. The atomic number tells you the element and, for a neutral atom, the number of electrons — and electrons determine chemical behavior. Isotope notation is the language of nuclear medicine (technetium-99m, iodine-131), radiocarbon dating (carbon-14), and nuclear power (uranium-235 vs. uranium-238), where a one-neutron difference can change a material's fate entirely. Average atomic masses are what you will multiply by moles in every stoichiometry calculation, so understanding why they are decimal (isotope-weighted averages, not single atoms) prevents a whole class of errors. The same isotope logic explains why iodine-131 is taken up by the thyroid gland like ordinary iodine — the chemistry is identical even though the mass differs.
The college version
Core Concepts
Atomic number: the element's identity card
The atomic number Z is the number of protons in the nucleus. It is the element's identity: no two elements share an atomic number, and changing Z changes the element. On the periodic table, Z is the whole number above the element symbol (C = 6, O = 8, Au = 79). In a neutral atom, the number of electrons also equals Z, because the total positive charge must balance the total negative charge.
Mass number and nuclide symbols
The mass number A is the count of protons plus neutrons. The number of neutrons is therefore:
N = A - Z
A specific atom with given Z and A is called a nuclide, written with the standard notation:
AZX
where X is the element symbol. For example, 23592U has 92 protons and 235 - 92 = 143 neutrons. Sometimes the notation omits the atomic number, since the element symbol already implies it (uranium-235, carbon-14, iodine-131 are all common spoken forms).
Isotopes: same element, different neutrons
Isotopes are atoms of the same element (same Z) with different numbers of neutrons (different A). Carbon has three naturally occurring isotopes: 126C, 136C, and 146C, with 6, 7, and 8 neutrons respectively. All three have 6 protons and 6 electrons (in neutral atoms), so their chemical behavior is essentially identical — but their masses differ, and some (like carbon-14) are radioactive. Isotopes are the reason average atomic masses on the periodic table are decimals.
Ions: unequal protons and electrons
When an atom gains or loses electrons, it becomes an Ion Atom with unequal proton and electron counts Full entry →. The charge is written as a superscript:
charge = Z - (number of electrons)
A neutral sodium atom (11 protons, 11 electrons) that loses one electron becomes Na+ with charge 11 - 10 = +1. A neutral chlorine atom (17 protons, 17 electrons) that gains one electron becomes Cl- with charge 17 - 18 = -1. Gaining or losing electrons does not change the atomic number — the element stays the same. Ions are the building blocks of ionic compounds, covered later in this chapter.
Average atomic mass: a weighted average
The atomic mass on the periodic table is not the mass of any single atom; it is the weighted average of all naturally occurring isotopes' masses, weighted by their Natural abundance The fraction of an element's atoms that are a given isotope Full entry →:
average atomic mass = ∑i ( fractional abundancei × massi )
where i runs over all isotopes. For example, chlorine is about 75.8% chlorine-35 (mass 34.97 u) and 24.2% chlorine-37 (mass 36.97 u). The weighted average comes out near 35.45 u — between the two isotopes, but pulled toward the more abundant one. The Atomic mass unit (u) Unit defined so carbon-12 = exactly 12 u Full entry → is defined so that one atom of carbon-12 has a mass of exactly 12 u.
Common Confusions
| Do Not Confuse | With | Difference |
|---|---|---|
| Atomic number (Z) | Mass number (A) | Z is protons only (defines element); A is protons + neutrons (defines the nuclide) |
| Mass number | Average atomic mass | Mass number counts particles in one specific nuclide (e.g., 235 for U-235); atomic mass is a weighted average over isotopes (e.g., 238.03 u for natural uranium) |
| Isotopes | Ions | Isotopes differ in neutrons (same element, same charge behavior); ions differ in electrons (same element, different charge). An atom can be both (e.g., 37Cl-) |
| Losing electrons | Losing protons | Losing electrons makes a positive ion but keeps the element; losing protons would change the element entirely — that only happens in nuclear reactions |
| Cation | Anion | Cation: positive, lost electrons (metals tend to form them). Anion: negative, gained electrons (nonmetals tend to form them). "Cat" is positive — a useful mnemonic |
| Abundance percentage | Isotope mass | The percentage tells you how common an isotope is (the weight in the average); the mass tells you how heavy it is (the value being weighted) |
Eli explains
The same idea, in plain words
Explain it like I’m 10
Every atom has an ID card. The number of protons on the card tells you what element it is — 6 protons always means carbon, like a fingerprint. Atoms of the same element can carry different numbers of neutrons, like twins wearing backpacks of different weights — same person, different total weight. We write the "weight class" on the card too, like 126C for a carbon with 12 total pieces. And when atoms gain or lose electrons, they get a plus or minus charge — like a balloon that gains or loses a few marbles. The decimal number on the periodic table is the average weight of all the twins in a big crowd, not any single one.
Worked example
Example 1 — Decoding a nuclide symbol. The nuclide 5626Fe (iron-56) is the most abundant isotope of iron. Identify the number of protons, neutrons, and electrons in a neutral atom, and write the symbol for its most common ion, Fe3+.
Read the symbol directly:
Z = 26, A = 56
Number of protons = Z = 26. Neutrons:
N = A - Z = 56 - 26 = 30
A neutral atom has electrons equal to protons, so 26 electrons. The Fe3+ ion has lost 3 electrons:
electrons in Fe3+ = 26 - 3 = 23
The ion still has 26 protons — the element is still iron.
Example 2 — Average atomic mass of chlorine. Chlorine consists of 75.78% chlorine-35 (isotope mass 34.97 u) and 24.22% chlorine-37 (isotope mass 36.97 u). Calculate the average atomic mass.
Use the weighted-average formula before substituting:
average atomic mass = ∑i ( fractional abundancei × massi )
Convert percentages to fractions and substitute:
= (0.7578 × 34.97 u) + (0.2422 × 36.97 u)
= 26.50 u + 8.954 u = 35.45 u
This matches the periodic-table value of 35.45 u. Notice the result is pulled toward 35 because the lighter isotope is far more abundant.
Example 3 — Which isotope has more neutrons? Compare the number of neutrons in carbon-14 and carbon-12, and explain why they are chemically nearly identical but differ in mass.
For each isotope, neutrons = A - Z, with Z = 6 for both:
N14C = 14 - 6 = 8 N12C = 12 - 6 = 6
Both have 6 protons and 6 electrons, so their electron configurations — and therefore their chemical reactions — are essentially the same. They differ only in neutron count, giving them different masses; carbon-14 is also radioactive, which is what makes it useful for dating.
Key takeaways
- Atomic number Z = number of protons; defines the element; in a neutral atom, electrons = Z.
- Mass number A = protons + neutrons; neutrons = A - Z.
- Nuclide notation: AZX; spoken form: element-name-mass number (uranium-235).
- Isotopes: same Z, different A (same protons, different neutrons); chemically near-identical, different masses, possibly radioactive.
- Ions: charge = Z - electrons; cations are positive (lost electrons), anions are negative (gained electrons); the element never changes.
- Average atomic mass = ∑(fractional abundance × isotope mass) — a weighted average, hence decimal values like 35.45 u for chlorine.
- Exam trap: the mass number is not the same as the atomic mass on the periodic table — one is a count for a specific nuclide, the other a weighted average over isotopes.
- Carbon-12 defines the atomic mass unit: exactly 12 u.
Check yourself
5 review questions from the chapter. Try each one, then open the answer.
How many protons, neutrons, and electrons are in a neutral atom of 4020Ca?
Show answer
Protons = 20 (Z), neutrons = 40 - 20 = 20, electrons = 20 (neutral atom).
Write the nuclide symbol for the isotope of oxygen with 10 neutrons.
Show answer
Oxygen has Z = 8. With 10 neutrons, A = 8 + 10 = 18, so the symbol is 188O (oxygen-18).
What is the charge on a sulfur ion with 16 protons and 18 electrons, and what kind of ion is it?
Show answer
Charge = 16 - 18 = -2, written S2-. It is an anion (gained 2 electrons).
Magnesium has isotopes 24Mg (78.99%, 23.985 u), 25Mg (10.00%, 24.986 u), and 26Mg (11.01%, 25.983 u). Calculate its average atomic mass.
Show answer
Average mass = (0.7899 × 23.985) + (0.1000 × 24.986) + (0.1101 × 25.983) = 18.946 + 2.499 + 2.861 = 24.31 u. This matches the periodic-table value for magnesium.
Why are isotopes of an element chemically almost identical even though they have different masses?
Show answer
Chemical behavior depends on the electron configuration, and all isotopes of an element have the same number of protons and electrons. Only the neutron count differs, which affects mass and nuclear stability, not electron arrangement.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- Atomic number (Z)
- The number of protons in a nucleus
- Mass number (A)
- Protons + neutrons in a nucleus
- Nuclide
- A specific atom with defined Z and A
- Isotope
- Atom of an element with a different neutron count
- Ion
- Atom with unequal proton and electron counts
- Cation / Anion
- Positive ion (lost electrons) / negative ion (gained electrons)
- Average atomic mass
- Abundance-weighted mean of isotope masses
- Atomic mass unit (u)
- Unit defined so carbon-12 = exactly 12 u
- Natural abundance
- The fraction of an element's atoms that are a given isotope
Sources & references
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