Subatomic Particles and Atomic Mass vs. Atomic Weight
An atom is built from three subatomic particles — protons, neutrons, and electrons — whose masses and abundances define atomic number, mass number, and atomic weight.
An atom is the smallest identifiable unit of an element: a neutral particle made of negatively charged electrons moving around a positively charged nucleus. Every atom is built from three subatomic particles — protons, neutrons, and electrons — and understanding how these particles are counted and weighed leads directly to two terms that are easy to confuse on the MCAT: mass number and atomic weight.
Key Takeaways
Protons (+1, ≈1 amu) and neutrons (0 charge, slightly >1 amu) sit in the nucleus; electrons (−1, ≈1/1,836 a proton's mass) orbit it, with valence electrons driving bonding.
Atomic number (Z) = number of protons (defines the element); mass number (A) = protons + neutrons; A − Z = neutrons. Mass number ≠ atomic weight.
Isotopes share atomic number but differ in mass number; atomic weight is the abundance-weighted average mass of all naturally occurring isotopes, in amu (1/12 the mass of carbon-12).
Moles connect the microscopic and macroscopic: 1 mole = Avogadro's number (6.022×10²³) of entities; molar mass links moles to grams.
Ordinary chemical reactions change only electron count/arrangement — proton/neutron counts change only in nuclear reactions.
An ion (cation if positive, anion if negative) forms when proton count ≠ electron count.
The Three Subatomic Particles
Protons sit in the nucleus and carry a charge equal to the fundamental unit of charge (≈1.6×10⁻¹⁹ C), denoted as a +1 charge. Each proton has a mass of approximately 1 atomic mass unit (amu).
Neutrons also sit in the nucleus but carry no charge. A neutron's mass is only slightly larger than a proton's, and together protons and neutrons account for almost the entire mass of an atom.
Electrons move through the space surrounding the nucleus at varying energy levels. Each electron carries a charge equal in magnitude to a proton's but opposite in sign (−1). An electron's mass is much smaller than a proton's — about 1/1,836 of it — which is why the electrostatic attraction between a proton and an electron dominates over the (comparatively tiny) gravitational attraction between their masses.
Particle | Charge | Relative Mass | Location |
|---|---|---|---|
Proton | +1 | ≈1 amu | Nucleus |
Neutron | 0 | Slightly more than 1 amu | Nucleus |
Electron | −1 | ≈1/1,836 of a proton | Orbiting the nucleus |
Electrons at greater distances from the nucleus occupy higher energy levels. The outermost electrons — the valence electrons — have the weakest attraction to the nucleus and the strongest interaction with the surrounding environment, which is why they're the electrons involved in chemical bonding.
Atomic Number and Mass Number
Two numbers define an atom's identity and composition:
Atomic number (Z): the number of protons in the nucleus. This determines the atom's chemical identity — every element on the periodic table has a unique atomic number.
Mass number (A): the total number of protons and neutrons (collectively called nucleons) in the nucleus.
Subtracting the atomic number from the mass number gives the number of neutrons: A − Z = neutrons. Mass number is not the same thing as atomic weight — that distinction is the focus of the next section.
Isotopes and Atomic Weight
Isotopes are variants of an element that share the same number of protons (same atomic number) but differ in their number of neutrons (different mass numbers). Isotopes matter beyond pure chemistry: stable isotopes are used in medical diagnostics, while radioactive isotopes are applied in treatment, archaeological dating, and as tracers in biochemical research.
Atomic weight is the weighted average of the masses of all naturally occurring isotopes of an element, measured in amu. Because most elements exist in nature as a mixture of isotopes, atomic weight isn't the mass of any single atom — it's an average that accounts for how abundant each isotope actually is.
Calculating atomic weight takes two inputs per isotope:
Isotope mass — the exact mass of that isotope (close to, but not exactly, its whole-number mass number, due to nuclear binding energy).
Natural abundance — the percentage of that isotope found in nature, expressed as a fraction.
The atomic weight is the sum of each isotope's mass multiplied by its natural abundance.
Worked example. Suppose a hypothetical element has two isotopes: Isotope A at 10 amu with 90% natural abundance, and Isotope B at 11 amu with 10% natural abundance. The atomic weight is: (10 amu × 0.90) + (11 amu × 0.10) = 9 + 1.1 = 10.1 amu.
This weighted average — 10.1 amu — is what would appear on a periodic table for this element, even though no single atom of it actually weighs 10.1 amu.
The Atomic Mass Unit (amu)
The atomic mass unit (amu) is defined as one-twelfth the mass of a carbon-12 atom (the stable isotope of carbon with six protons and six neutrons), which makes it a convenient, universal standard for expressing atomic-scale masses. One amu is approximately 1.66×10⁻²⁷ kg. The atomic mass listed for an element on the periodic table is expressed in amu and represents that element's atomic weight — the abundance-weighted average across its naturally occurring isotopes.
Connecting Mass, amu, and Moles
A mole is a unit for counting amounts of substance, the same way "a dozen" signals 12 of something. One mole of any substance contains exactly 6.022×10²³ entities (atoms, molecules, ions) — this is Avogadro's number. The mole lets chemists count particles by weighing them, since the mass of one mole of a substance equals its atomic or molecular mass expressed in grams — its molar mass (g/mol).
These three quantities — number of particles, moles, and mass — convert into each other:
Particles ↔ moles: divide by Avogadro's number to get moles; multiply by Avogadro's number to get particles.
Moles ↔ mass: multiply moles by molar mass to get mass; divide mass by molar mass to get moles.
Particles ↔ mass: chain through moles — divide by molar mass then multiply by Avogadro's number (particles from mass), or divide by Avogadro's number then multiply by molar mass (mass from particles).
These conversions bridge the microscopic world (atoms, amu) and the macroscopic world (grams, a measurable mass) — essential for balancing chemical equations and scaling reactions.
One more distinction matters here: in ordinary chemical reactions, the nucleus never changes — only the number and arrangement of electrons can change. Proton or neutron counts change only during nuclear reactions. This is why the movement of electrons, not protons or neutrons, is the focus of chemical bonding and reactivity.
Ions: Cations and Anions
An ion is a charged particle that results when an atom's proton count no longer equals its electron count. A cation is a positively charged ion (more protons than electrons); an anion is a negatively charged ion (more electrons than protons). Because ions carry a net charge, their chemical properties differ meaningfully from those of the neutral atom they came from — which is why it matters, throughout chemistry, whether a species is written as a neutral atom or a charged ion.
Common MCAT Mistakes
Confusing mass number with atomic weight. Mass number (A) is a whole-number count of protons plus neutrons in one specific atom; atomic weight is the decimal, abundance-weighted average across all of an element's naturally occurring isotopes.
Forgetting that neutron count doesn't change an atom's identity. Isotopes of the same element share the same atomic number (proton count) but differ in mass number (neutron count) — the element's identity is set by protons alone.
Treating electron count as always equal to proton count. That's only true for a neutral atom. An ion forms whenever proton count and electron count differ — a cation has fewer electrons than protons, an anion has more.
Treating a mole as a mass. A mole is a count of entities (6.022×10²³ of them), not a mass. The mass of one mole of a substance equals its molar mass in grams — the two quantities are linked by molar mass, not identical to each other.
MCAT-Style Concept Check
Question: An element has two naturally occurring isotopes: Isotope X at 63 amu with 70% natural abundance, and Isotope Y at 65 amu with 30% natural abundance. What is this element's atomic weight, as it would appear on the periodic table?
A) 63.0 amu
B) 63.6 amu
C) 64.0 amu
D) 65.0 amu
Answer: B
Explanation: Atomic weight is the sum of each isotope's mass multiplied by its natural abundance: (63 amu × 0.70) + (65 amu × 0.30) = 44.1 + 19.5 = 63.6 amu. This weighted average sits closer to 63 amu than 65 amu because Isotope X is more abundant — it's not a simple midpoint between the two isotope masses.
FAQ
What is the difference between mass number and atomic weight?
Mass number is the whole-number total of protons and neutrons in a single atom's nucleus. Atomic weight is the abundance-weighted average mass, in amu, of all of an element's naturally occurring isotopes — which is why atomic weight on the periodic table is almost always a decimal, not a whole number.
What are isotopes?
Isotopes are atoms of the same element — same atomic number, same number of protons — that have different numbers of neutrons, and therefore different mass numbers. Because they contain the same number of protons, isotopes of an element share the same chemical identity.
What is a mole, and why does Avogadro's number matter?
A mole is a counting unit, like "a dozen," equal to exactly 6.022×10²³ entities — this is Avogadro's number. It matters because it lets chemists convert between a measurable mass (grams) and a specific number of atoms or molecules, using the substance's molar mass as the bridge.
What's the difference between a cation and an anion?
Both are ions — atoms or molecules with a net charge because their proton count no longer equals their electron count. A cation has more protons than electrons, giving it a net positive charge; an anion has more electrons than protons, giving it a net negative charge.
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