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General Chemistry I · Atomic Structure

Atomic Theory and Structure

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On this page 7 sections
  1. In 30 seconds
  2. Why this matters
  3. The college version
  4. Eli explains
  5. Worked example
  6. Key takeaway
  7. Study tools

In 30 seconds

Atoms consist of a tiny, dense, positively charged (protons and neutrons) surrounded by a cloud of electrons. The counts protons and defines the element; the is protons plus neutrons. Isotopes are atoms of the same element with different numbers of neutrons. is a weighted average of masses based on natural abundance.

Why this matters

Isotopes matter in medicine: radioactive iodine-131 is used to treat thyroid conditions because the thyroid gland concentrates iodine. Carbon-14 dating uses the C-14/C-12 ratio to estimate the age of organic artifacts. In mass spectrometry, a molecule's isotope pattern acts as a fingerprint for identification—a core tool in toxicology and drug testing.

The college version

1. Discovery of Atomic Structure

  • John Dalton (1803): proposed atoms as indivisible spheres, reviving the atomic theory.
  • J.J. Thomson (1897): using a cathode-ray tube, discovered the —a negatively charged particle far lighter than any atom. He proposed the "plum pudding" model: negative electrons embedded in a positive sphere.
  • Robert Millikan (1909): his oil-drop experiment measured the electron's charge and mass precisely (charge = −1.602 × 10⁻¹⁹ C).
  • Ernest Rutherford (1911): in the , most alpha particles passed straight through a thin gold foil, but a few bounced back—revealing a tiny, dense, positively charged nucleus. He proposed the nuclear model.

2. Subatomic Particles

ParticleSymbolChargeApprox. mass (amu)Location
Protonp⁺+11.0073nucleus
Neutronn⁰01.0087nucleus
Electrone⁻−10.00055outside nucleus

Protons and neutrons make up essentially all of an atom's mass; electrons are about 1/1836 the mass of a but occupy the atom's volume.

3. Atomic Number, Mass Number, and Isotopes

  • Atomic number (Z): the number of protons; it defines the element. In a neutral atom, protons = electrons.
  • Mass number (A): protons + neutrons.
  • Isotopes: atoms with the same Z but different A (different counts). Notation: AZX or X-A (e.g., carbon-14, 146C).
  • Ions: atoms that have gained or lost electrons; charge = protons − electrons.

How it works

  1. Thomson's cathode-ray tube showed a negative particle (the electron) lighter than any atom.
  2. Millikan's oil-drop experiment measured the electron's charge and mass.
  3. Rutherford's gold-foil experiment showed most of the atom is empty and its mass sits in a tiny positive nucleus.
  4. Chadwick (1932) discovered the neutron, explaining the nucleus's extra mass.
  5. Count particles: protons = Z; neutrons = A − Z; electrons = protons (neutral) or adjusted for ionic charge.
  6. Weight isotope masses by natural abundance to obtain average atomic mass.

Common confusions

Do not confuseWithDifference
Atomic numberMass numberZ = protons; A = protons + neutrons
IsotopeIonIsotope = neutron count varies; ion = electron count varies
Atomic massMass numberAtomic mass = weighted average (amu); mass number = integer count
NucleusAtomNucleus is the tiny core; the atom includes the electron cloud

Memory aids

"Zero in on Z" — the atomic number Z is the proton count, the element's ZIP code that pinpoints its identity. "A = All nucleons" — the mass number A adds up all the heavy particles (protons + neutrons).

Quick review

Topic Recap

Experiments by Thomson, Millikan, and Rutherford revealed the electron, its charge/mass, and the dense positive nucleus. Atoms are defined by atomic number (protons) and mass number (protons + neutrons). Isotopes vary in neutrons, and average atomic mass is the abundance-weighted average of isotope masses.

Knowledge Check

  1. An atom has 17 protons and 20 neutrons. What are Z, A, and the number of electrons (neutral)?
  2. Which scientist's experiment revealed the nucleus?
  3. Boron has two isotopes: B-10 (10.013 , 19.9%) and B-11 (11.009 amu, 80.1%). What is the average atomic mass?
  4. How many neutrons are in uranium-235 (Z = 92)?
  5. What distinguishes isotopes of the same element?

Answers and Rationales

  1. Z = 17, A = 37, electrons = 17 (neutral). This is chlorine-37.
  2. Rutherford — the gold-foil experiment showed a dense positive nucleus.
  3. 10.81 amu — (0.199 × 10.013) + (0.801 × 11.009) = 10.81.
  4. 143 — neutrons = A − Z = 235 − 92 = 143.
  5. Different numbers of neutrons (same number of protons/atomic number).
Eli, the EliExplains learning guide

Eli explains

The same idea, in plain words

Explain it like I’m 10

Picture a football stadium. The nucleus is a marble sitting at the 50-yard line; the electrons are tiny gnats buzzing throughout the entire stadium. Most of the atom is empty space. The marble (nucleus) is heavy and positively charged; the electrons are light and negatively charged, attracted to the positive nucleus but never crashing into it.

Where it stops being exact: electrons do not orbit in neat circles like tiny planets. Modern quantum mechanics describes them as clouds of probability (covered in later topics), and the "empty space" is really filled by the electron cloud and the electric field, not true nothingness.

Simple Example

Carbon-12 has 6 protons, 6 neutrons, and 6 electrons. Carbon-14 (used in radiocarbon dating) also has 6 protons but 8 neutrons—the same element, a different isotope.

Worked example

Average Atomic Mass

Average atomic mass = ∑(fractional abundance × isotope mass)

Example 1. Chlorine has two naturally occurring isotopes: Cl-35 (34.969 amu, 75.77% abundance) and Cl-37 (36.966 amu, 24.23% abundance). Find the average atomic mass. (0.7577 × 34.969 amu) + (0.2423 × 36.966 amu) = 26.50 + 8.956 = 35.45 amu This matches the periodic-table value for chlorine.

Example 2. An element has two isotopes: X-10 (10.013 amu, 19.9%) and X-11 (11.009 amu, 80.1%). Find the average atomic mass. (0.199 × 10.013) + (0.801 × 11.009) = 1.993 + 8.818 = 10.81 amu The element is boron.

Common errors: using percentages without converting to decimals (divide by 100); omitting one of the isotopes; and confusing the mass number (an integer count) with atomic mass (a weighted average in amu).

Key takeaways

  • High yield: Atomic number = protons = element identity; changing protons changes the element.
  • High yield: Mass number A = protons + neutrons; isotopes differ only in neutron count.
  • High yield: Average atomic mass is a weighted average, not a simple average.
  • High yield: Neutral atom: electrons = protons; ion: electrons = protons − charge.
  • Rutherford: tiny dense positive nucleus; Thomson: electron; Millikan: electron charge/mass.
  • 1 amu ≈ the mass of one proton or one neutron.

Keep learning

Ready to build on this? Continue to the next lesson.

Practice General Chemistry I

This lesson has no separate scored set. Practice draws from the subject’s question bank.

Study tools & related lessonsYou’ll learn to · Key vocabulary · Related

You’ll learn to

  • Describe how the experiments of Thomson, Millikan, and Rutherford revealed atomic structure.
  • List the three subatomic particles with their charges, relative masses, and locations.
  • Use atomic number and mass number to determine the numbers of protons, neutrons, and electrons.
  • Calculate average atomic mass from isotope masses and natural abundances.

Key vocabulary

Cathode ray
Stream of electrons
Plum pudding model
Electrons in a positive sphere
Gold-foil experiment
Alpha-particle scattering
Nucleus
Dense positive core
Proton
+1 particle in the nucleus
Neutron
Neutral nuclear particle
Electron
−1 particle outside the nucleus
Atomic number (Z)
Number of protons
Mass number (A)
Protons + neutrons
Isotope
Same Z, different A
Ion
Atom with a net charge
amu
Atomic mass unit
Average atomic mass
Weighted isotope average

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