Chemistry 2e · Atoms, Molecules, and Ions
Evolution of Atomic Theory
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In 30 seconds
Dalton's atoms were indivisible — but they did not stay that way for long. Between 1897 and 1932, a sequence of ingenious experiments revealed that the atom has an internal structure made of electrons, protons, and neutrons. J. J. Thomson discovered the Electron Negatively charged subatomic particle with charge −1 and mass 9.11 × 10-28 g Full entry → in cathode rays and proposed the "plum pudding" model; Robert Millikan measured the electron's charge with oil droplets; Ernest Rutherford's gold-foil experiment shattered the plum pudding and revealed a tiny, dense, positively charged Nucleus Tiny, dense, positively charged center of the atom containing protons and neutrons Full entry →; and James Chadwick finally identified the Neutron Neutral particle in the nucleus, mass nearly equal to the proton's Full entry →. Each model replaced the previous one because new measurements demanded it — the perfect case study of how science actually advances: not by clever arguments, but by experiments that discriminate between competing models.
Why this matters
The subatomic particles discovered in this era are the vocabulary of all modern chemistry. The Proton Positively charged particle in the nucleus; charge +1, mass 1.67 × 10-24 g Full entry → count defines the element (atomic number); the electron arrangement determines chemical behavior (bonding, reactivity, redox); the neutron count distinguishes isotopes (which matters for nuclear medicine, radiocarbon dating, and reactor design). Understanding how these particles were discovered also explains what the atom really is: mostly empty space, with nearly all its mass in a nucleus roughly 100,000 times smaller than the atom itself. That counterintuitive fact drives technologies from MRI magnets to particle accelerators, and it explains why alpha-emitting materials are hazardous — fast, massive, charged particles tear through matter. The experiment–model–refinement loop shown here is the template for every theory in this book.
The college version
Core Concepts
Cathode rays and the discovery of the electron
In the 1890s, physicists studied cathode rays — glowing beams produced in evacuated glass tubes with a high voltage applied across two metal electrodes. Experiments showed the rays were deflected by magnetic and electric fields, which meant they carried electric charge. By measuring how much the rays bent, J. J. Thomson (1897) showed they were streams of negatively charged particles with a specific charge-to-mass ratio (e/m). Astonishingly, the same e/m ratio was obtained no matter what metal formed the electrodes or what gas filled the tube — the particles were a universal constituent of all matter. Thomson had discovered the electron, the first subatomic particle. His measured charge-to-mass ratio, about 1.76 × 108 C/kg, is consistent with modern measurements.
Millikan's oil-drop experiment: the electron's charge
Knowing the electron's charge-to-mass ratio but not its charge, you can't find its mass. Robert Millikan solved this in 1909 with the oil-drop experiment: he sprayed fine oil droplets between charged plates, watched them fall under gravity, and adjusted the voltage until individual droplets hung motionless, balanced between electrical and gravitational forces. Repeating this for many droplets, he found every charge was a whole-number multiple of a single fundamental value — the charge of one electron:
e = 1.602 × 10-19 C
Combining that with Thomson's e/m gave the electron's mass: about 9.11 × 10-28 g — roughly 1/1836 the mass of the lightest atom, hydrogen.
The plum pudding model
To accommodate negatively charged electrons inside neutral atoms, Thomson proposed the Plum pudding model Thomson's model: diffuse positive sphere with embedded electrons Full entry →: a diffuse sphere of positive charge (the "pudding") with electrons ("plums") embedded throughout, like raisins in a cake. It was reasonable and explained why atoms are electrically neutral, but it made a specific prediction about how alpha particles would pass through matter — and that prediction failed.
Radioactivity: a new probe
Henri Becquerel (1896) discovered that uranium salts emit penetrating rays spontaneously, and Marie and Pierre Curie isolated new radioactive elements (polonium, radium). Ernest Rutherford identified three types of emission: alpha particles (positive, helium nuclei), beta particles (negative, electrons), and gamma rays (high-energy electromagnetic radiation). Alpha particles became the perfect probe: fast, heavy, and positively charged, they could be fired at thin metal foils to map what was inside atoms.
Rutherford's gold-foil experiment: the nucleus
In 1909–1911, Hans Geiger and Ernest Marsden, working in Rutherford's lab, fired alpha particles at an extremely thin gold foil. The plum pudding model predicted the alphas would pass through with only slight deflections. Instead, most passed straight through, but a small fraction deflected at large angles — some even backward. Rutherford famously said it was "as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you." The only explanation: nearly all of an atom's mass and all of its positive charge are concentrated in a tiny central nucleus, with electrons orbiting the mostly empty space around it. The nuclear model was born.
The discovery of the neutron
The nuclear model had a problem: helium's nucleus was measured at 4 atomic mass units but contained only 2 units of positive charge. Where was the extra mass? In 1932, James Chadwick bombarded beryllium with alpha particles and detected a new, electrically neutral radiation — the neutron, with mass nearly identical to the proton's. The neutron explained nuclear masses and made isotope chemistry possible. Its discovery also highlighted an experimental lesson: neutral particles are invisible to electric and magnetic fields, so they are much harder to detect than charged ones.
The modern picture
| Particle | Charge | Approximate mass (g) | Location |
|---|---|---|---|
| Electron (e-) | −1 | 9.11 × 10-28 | Outside the nucleus |
| Proton (p+) | +1 | 1.67 × 10-24 | In the nucleus |
| Neutron (n0) | 0 | 1.67 × 10-24 | In the nucleus |
A proton is about 1836 times more massive than an electron, so the nucleus carries essentially all of an atom's mass. In an electrically neutral atom, the number of electrons equals the number of protons. These identities feed directly into the next topic, where they become atomic numbers, mass numbers, and isotope symbols.
Common Confusions
| Do Not Confuse | With | Difference |
|---|---|---|
| The electron's discovery | The electron's charge measurement | Thomson found electrons and measured e/m; Millikan measured the absolute charge e years later. Both were needed for the mass |
| Alpha particles | Beta particles | Alpha = positive, heavy (helium nuclei), stopped by paper; beta = negative, light (electrons), stopped by thin metal. Their behavior in fields and matter differs completely |
| Plum pudding model | Nuclear model | Plum pudding: positive charge spread throughout, electrons embedded. Nuclear: all positive charge and nearly all mass in a tiny central nucleus. The gold-foil experiment decided between them |
| An atom's size | Its nucleus's size | The atom (electron cloud) is ~10⁻¹⁰ m across; the nucleus is ~10⁻¹⁵ m — about 100,000 times smaller, yet holds ~99.97% of the mass |
| Neutron mass | Proton mass | Nearly equal (both ~1.67 × 10⁻²⁴ g), but the neutron is slightly heavier — a small difference with big consequences in nuclear physics |
| "Indivisible" Dalton atom | Modern atom | Dalton's atoms were fundamental particles; today atoms are known to be made of protons, neutrons, and electrons — a model change, not a rejection of atomic theory |
Eli explains
The same idea, in plain words
Explain it like I’m 10
Scientists used to think atoms were like tiny, solid marbles that couldn't be broken. Then they found electricity shooting through glass tubes and discovered a tiny piece inside the marble — the electron — that is the same in everything, like finding the same Lego brick inside every toy. A few years later, they shot tiny bullets at gold foil and were shocked when some bounced back, which showed that an atom is mostly empty space with a super-heavy, tiny center called the nucleus. It's like discovering a building is 99.99% air, with all the weight in a single doorknob in the middle.
Worked example
Example 1 — From charge to mass. Thomson measured the electron's charge-to-mass ratio as 1.76 × 108 C/kg, and Millikan measured the charge as 1.602 × 10-19 C. Find the electron's mass.
The charge-to-mass ratio is defined as:
em = 1.76 × 108 Ckg
Rearrange to solve for mass, then substitute:
m = ee/m = 1.602 × 10-19 C1.76 × 108 C/kg = 9.10 × 10-28 kg
Convert to grams: 9.10 × 10-28 kg × 1000 g1 kg = 9.10 × 10-25 g. The modern accepted value is 9.109 × 10-28 g — the agreement is excellent.
Example 2 — Relative masses. The proton's mass is 1.6726 × 10-24 g and the electron's is 9.109 × 10-28 g. How many times more massive is the proton?
Set up the ratio:
mprotonmelectron = 1.6726 × 10-24 g9.109 × 10-28 g ≈ 1836
A proton is about 1836 times heavier than an electron. This is why electron mass is often ignored in atomic mass calculations — it is negligible next to protons and neutrons.
Example 3 — Charges add up. A neutral carbon atom has 6 protons and 6 electrons. What is the total charge, and what happens if one electron is removed?
Charge is the sum of proton and electron charges:
Q = 6(+1) + 6(-1) = 0
Removing one electron leaves:
Q = 6(+1) + 5(-1) = +1
The result is a positively charged ion — the same logic that later becomes oxidation states and ionic bonding.
Key takeaways
- Thomson (1897): discovered the electron using cathode rays; measured e/m ≈ 1.76 × 108 C/kg; proposed the plum pudding model.
- Millikan (1909): oil-drop experiment measured the electron charge e = 1.602 × 10-19 C; combining with Thomson's ratio gives electron mass 9.11 × 10-28 g.
- Rutherford (1911): gold-foil experiment — most alpha particles pass through, a few deflect sharply, some bounce back → the atom is mostly empty space with a tiny, dense, positive nucleus.
- Chadwick (1932): discovered the neutron, a neutral particle with mass nearly equal to the proton's.
- Radioactivity: alpha (+) = helium nuclei, beta (−) = electrons, gamma = high-energy photons.
- Proton ≈ 1836 × electron mass; the nucleus holds ~99.97% of an atom's mass in ~1/100,000 of its volume.
- In a neutral atom: number of electrons = number of protons.
- Safety principle (general, not a procedure): radioactive materials are handled with shielding, distance, and minimal exposure time whenever sources are used.
Check yourself
5 review questions from the chapter. Try each one, then open the answer.
What experimental evidence forced physicists to abandon the plum pudding model?
Show answer
Rutherford's gold-foil experiment: most alpha particles passed straight through the foil, but a few deflected sharply and some bounced back. A diffuse positive sphere (plum pudding) could not produce such large deflections; only a tiny, dense, positive nucleus could.
Why couldn't Thomson's e/m measurement alone give the electron's mass?
Show answer
Because e/m is a ratio — two unknowns. Millikan's independent measurement of e supplied the second equation, letting Thomson's ratio be converted into a mass.
What did Millikan's oil-drop experiment actually measure, and what was the key result?
Show answer
The charge of a single electron, 1.602 × 10-19 C. The key result: every droplet's charge was an integer multiple of that value, showing charge comes in discrete packets.
How did Chadwick's neutron solve the "missing mass" problem of the nucleus?
Show answer
Helium's nucleus has 2 protons but measured 4 atomic mass units; protons alone couldn't supply the mass. Chadwick found a neutral particle of nearly proton mass, and nuclei turned out to contain neutrons as well as protons.
Roughly what fraction of an atom's volume is occupied by its nucleus, and why does that matter for how radiation passes through matter?
Show answer
The nucleus is about 1/100,000 of the atom's diameter, so it occupies roughly 10⁻¹⁵ of the volume. That's why most alpha particles pass through gold foil untouched — matter is mostly empty space — while the rare direct hit on the nucleus causes a dramatic bounce.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- Cathode ray
- Beam of electrons produced in an evacuated tube under high voltage
- Electron
- Negatively charged subatomic particle with charge −1 and mass 9.11 × 10-28 g
- Proton
- Positively charged particle in the nucleus; charge +1, mass 1.67 × 10-24 g
- Neutron
- Neutral particle in the nucleus, mass nearly equal to the proton's
- Nucleus
- Tiny, dense, positively charged center of the atom containing protons and neutrons
- Plum pudding model
- Thomson's model: diffuse positive sphere with embedded electrons
- Alpha particle
- Helium nucleus (42He2+) emitted in radioactive decay
- Charge-to-mass ratio (e/m)
- The electron's charge divided by its mass
Sources & references
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