2.2 Discovery of Atomic Structure: Cathode Rays and the Nuclear Model

For over a century, Dalton's atomic theory treated the atom as the ultimate indivisible particle. This section is where that idea collapses and modern chemistry begins. If you understand how electrons were discovered, why radioactivity forced scientists to rethink the atom, and how Rutherford proved the atom is mostly empty space with a dense nucleus, you unlock the entire logic of periodic trends, bonding, and nuclear chemistry that dominates A-Level, AP, IB and College exams.

Key Insights

• Cathode rays are streams of negatively charged electrons originating at the cathode in an evacuated tube.
• J.J. Thomson measured the electron charge-to-mass ratio and Millikan measured its charge to calculate its mass.
• Radioactivity consists of alpha particles, beta particles which are electrons, and neutral gamma radiation.
• Thomson's plum-pudding model proposed electrons embedded in a uniform positive sphere.
• Rutherford's gold foil experiment proved the nuclear model with a tiny dense positive nucleus and mostly empty space.

The Discovery of Atomic Structure

Dalton based his atomic theory on chemical observations made in the laboratory. Neither he nor those who followed him during the century after his work was published had any direct evidence for the existence of atoms. Atoms were a useful explanatory model for the law of definite proportions and the law of multiple proportions, but not an observed reality.

Today, that has changed completely. With scanning tunneling microscopy, we can measure the properties of individual atoms and even provide images of them. The surface of silicon, for example, appears as an ordered array of spheres where each purple sphere is a single silicon atom. This capability proves Dalton's hypothesis was physically correct.

As scientists developed methods for probing the nature of matter, the supposedly indivisible atom began to show signs of a more complex structure, and today we know that the atom is composed of subatomic particles. Before we summarize the current model, we briefly consider a few of the landmark discoveries that led to that model. We will see that the atom is composed in part of electrically charged particles, some with a positive charge and some with a negative charge. As we discuss the development of our current model of the atom, keep in mind this fact: Particles with the same charge repel one another, whereas particles with unlike charges attract one another.

Callout Box: The Big Shift> Dalton's atom was solid, indivisible, and electrically neutral as a whole. The discovery of subatomic particles meant atoms have internal structure, are made of charged components, and can be broken apart in experiments. This shift explains why atoms interact, bond, and undergo nuclear changes.

Cathode Rays and Electrons

During the mid-1800s, scientists began to study electrical discharge through a glass tube pumped almost empty of air. This apparatus is called a cathode-ray tube. When a high voltage was applied to the electrodes in the tube, radiation was produced between the electrodes. This radiation, called cathode rays, originated at the negative electrode, called the cathode, and traveled to the positive electrode, called the anode. Although the rays could not be seen, their presence was detected because they cause certain materials to fluoresce, or to give off light.

Two simple tests revealed the direction and charge of cathode rays:

1. Proof of direction: A small glass screen was placed diagonally inside the tube. The cathode rays struck the screen and produced a bright fluorescent line, showing the path from cathode (-) to anode (+). This proved travel from cathode to anode.
2. Proof of charge: When a magnet was brought near the tube, the fluorescent path was deflected. Since moving charged particles are deflected by magnetic fields, this showed cathode rays are charged.

Photorealistic 3D render of J.J. Thomson cathode-ray tube with perpendicular electric and magnetic fields, showing cathode, anode, electrically charged plates, magnet, electron path, fluorescent screen

Experiments showed that cathode rays are deflected by electric or magnetic fields in a way consistent with their being a stream of negative electrical charge. The British scientist J. J. Thomson (1856-1940) observed that cathode rays are the same regardless of the identity of the cathode material. In a paper published in 1897, Thomson described cathode rays as streams of negatively charged particles. This paper is generally accepted as the discovery of what became known as the electron.

Thomson's Charge-to-Mass Experiment

Thomson constructed a cathode-ray tube having a hole in the anode through which a beam of electrons passed. Electrically charged plates and a magnet were positioned perpendicular to the electron beam, and a fluorescent screen was located at one end. The electric field deflected the rays in one direction, and the magnetic field deflected them in the opposite direction.

Thomson adjusted the strengths of the electric and magnetic fields so that the effects balanced each other, allowing the electrons to travel in a straight path to the screen. When the forces cancel:

Electric force = Magnetic force
\(eE\) = \(e v B\)

Knowing the strengths that resulted in the straight path made it possible to determine the velocity of the electrons, and from the deflection when only one field was applied, he could calculate the ratio of the electron's electrical charge to its mass:

Charge-to-mass ratio = 1.76 × 10⁸ coulombs per gram (C/g)

This value is enormous. It means the electron has a very small mass for its charge. This ratio was the same regardless of the gas in the tube or the metal used for electrodes, proving electrons are a universal component of all matter.

Pro Tip for Exams:You must know the direction of deflection. In an electric field, electrons (negative) are attracted to the positive plate and repelled by the negative plate. If no magnetic field were applied, the electron beam would be deflected toward the positive electrically charged plate. Examiners love this.

Millikan's Oil-Drop Experiment - Measuring the Electron's Charge and Mass

Once the charge-to-mass ratio of the electron was known, measuring either the charge or the mass would allow calculation of the other. In 1909, Robert Millikan (1868-1953) of the University of Chicago succeeded in measuring the charge of an electron by performing the oil-drop experiment.

Photorealistic 3D cutaway of Millikan oil-drop apparatus showing oil drops falling through hole, electrically charged plates, X-ray source, and viewing microscope

How it worked:

1. Small drops of oil were sprayed into a chamber above two electrically charged plates. The drops fell through a small hole in the upper plate.
2. The drops were ionized by X-rays, picking up extra electrons and becoming negatively charged.
3. Between the electrically charged plates, Millikan varied the voltage. By adjusting the voltage, he could make the negatively charged oil drops hang suspended in air, balancing gravitational attraction against electrical repulsion.
4. He measured the rate of fall of the drops and calculated the negative charge on each drop.
5. Because the charge on any drop was always some integral multiple of 1.602 × 10^-19 C, Millikan deduced this value to be the charge of a single electron.

He then calculated the mass of the electron by using his experimental value for the charge and Thomson's charge-to-mass ratio:

Electron mass = Charge / (Charge/mass ratio)
Electron mass = (1.602 × 10^-19 C) / (1.76 × 10⁸ C/g) = 9.10 × 10^-28 g

This result agrees well with the currently accepted value of 9.10938 × 10^-28 g. This mass is about 2000 times smaller than that of hydrogen, the lightest atom.

Property	Thomson's Finding	Millikan's Finding
What was measured	Charge-to-mass ratio (e/m)	Charge of single electron (e)
Value	1.76 × 108 C/g	1.602 × 10-19 C
Significance	Proved electrons are fundamental particles with tiny mass	Allowed first calculation of electron mass

Exam Trap:Students confuse C/g with g/C and invert the calculation. Remember: Thomson gave C per gram, so you divide Coulombs by C/g to get grams. Also, do not write electron charge as -1.602 × 10^-19 C in the mass calculation - use magnitude. The negative sign indicates charge type, not magnitude for mass.

Radioactivity

In 1896 the French scientist Henri Becquerel (1852-1908) discovered that a compound of uranium spontaneously emits high-energy radiation. This spontaneous emission of radiation is called radioactivity. At Becquerel's suggestion, Marie Sklodowska Curie (1867-1934) and her husband, Pierre, began experiments to isolate the radioactive components of the compound. When Marie Curie presented her doctoral thesis, it was described as the greatest single contribution of any doctoral thesis in the history of science. In 1903 Becquerel, Marie Curie, and her husband were jointly awarded the Nobel Prize in Physics for their pioneering work on radioactivity, a term she introduced. In 1911 Marie Curie won a second Nobel Prize, this time in chemistry, for her discovery of the elements polonium and radium.

Further study of radioactivity, principally by the British scientist Ernest Rutherford (1871-1937), revealed three types of radiation: alpha (α), beta (β), and gamma (γ). The paths of alpha and beta radiation are bent by an electric field, although in opposite directions; gamma radiation is unaffected by the field.

Photorealistic 3D render of alpha beta gamma radiation behavior in electric field with lead block, radioactive source, electrically charged plates, and photographic plate

Rutherford showed that alpha and beta rays consist of fast-moving particles. In fact:

• Beta particles are high-speed electrons and can be considered the radioactive equivalent of cathode rays. They are attracted to a positively charged plate. Charge = 1- in units of electron charge.
• Alpha particles have a positive charge and are attracted to a negative plate. Charge = 2+. Each alpha particle has a mass about 7400 times that of an electron. Today we know an alpha particle is a helium nucleus, He²⁺.
• Gamma radiation is high-energy radiation similar to X-rays; it does not consist of particles and carries no charge. It is not deflected.

Why are beta rays deflected to a greater extent than alpha rays? Because deflection depends on charge-to-mass ratio. Beta particles have a much larger charge-to-mass ratio and much smaller mass, so the same electric force produces a larger acceleration.

Type of Radiation	Symbol	Charge	Mass (relative to electron)	Identity	Deflection in Electric Field
Alpha	α	2+	∼7400 × electron	Helium nucleus	Toward negative plate, small deflection
Beta	β	1-	1 × electron	High-speed electron	Toward positive plate, large deflection
Gamma	γ	0	0 (no particle)	High-energy electromagnetic radiation	No deflection

Callout Box:Why Radioactivity Mattered > Radioactivity provided the first evidence that atoms are not immutable. If atoms emit particles with different masses and charges, they must contain those particles. It also gave Rutherford a perfect projectile - the massive, positively charged alpha particle - for probing the interior of the atom.

Exam Tip:In A-Level and AP, you must state: beta is an electron from the nucleus, not from electron shells. Its origin is neutron decay. Alpha is a helium nucleus. Gamma is electromagnetic wave with no charge or mass. Many students write gamma has charge - instant loss.

The Nuclear Model of the Atom

With growing evidence that the atom is composed of smaller particles, attention was given to how the particles fit together. During the early 1900s, Thomson reasoned that because electrons contribute only a very small fraction of the atom's mass, they probably are responsible for an equally small fraction of the atom's size. He proposed that the atom consisted of a uniform positive sphere of matter in which the electrons were embedded like raisins in a pudding or seeds in a watermelon. This plum-pudding model, named after a traditional English dessert, was very short-lived.

In 1910, Rutherford was studying the angles at which alpha particles were deflected, or scattered, as they passed through a thin sheet of gold foil. He discovered that almost all the particles passed directly through the foil without deflection, with a few particles deflected about 1 degree, consistent with Thomson's plum-pudding model. For the sake of completeness, Rutherford suggested that Ernest Marsden, an undergraduate student working in the laboratory, look for scattering at large angles. To everyone's surprise, a small amount of scattering was observed at large angles, with some particles scattered back in the direction from which they had come. The explanation for these results was not immediately obvious, but they were clearly inconsistent with Thomson's plum-pudding model.

Photorealistic 3D render of Rutherford gold foil alpha scattering experiment showing source, beam, gold foil, circular fluorescent screen, and concept of electrostatic repulsion between positive alpha particle and positive nucleus

The Rutherford Scattering Experiment - Setup

• Source of alpha particles housed in a lead block with a small opening to produce a narrow beam.
• Beam directed at an ultrathin sheet of gold foil, only a few atoms thick.
• Surrounding the foil, a circular fluorescent screen that scintillates when struck by alpha particles, allowing detection at all angles.

Observations:

1. Most alpha particles underwent no scattering because most of the atom is empty. They passed straight through.
2. A few were deflected by about 1 degree.
3. About 1 in 20,000 was deflected at large angles >90 degrees, even bouncing back.

According to the plum-pudding model, the positive charge is spread out, so the electric field would be weak and could cause only very minor deflections. Large-angle scattering should be impossible.

Rutherford's Nuclear Model

Rutherford explained the results by postulating the nuclear model of the atom, a model in which most of the mass of each atom and all of its positive charge reside in a very small, extremely dense region that he called the nucleus. He postulated further that most of the volume of an atom is empty space in which electrons move around the nucleus.

In the alpha-scattering experiment, most particles passed through foil unscattered because they did not encounter the minute nucleus of any gold atom. Occasionally, however, an alpha particle came close to a gold nucleus. The repulsion between the highly positive charge of the gold nucleus and the positive charge of the alpha particle was strong enough to deflect the particle, with a head-on collision causing backward scattering.

Feature	Thomson Plum-Pudding Model	Rutherford Nuclear Model
Positive charge	Spread uniformly throughout sphere	Concentrated in tiny central nucleus
Electrons	Embedded in positive sphere like raisins	Orbiting around nucleus in empty space
Mass distribution	Uniform	Almost all mass in nucleus
Predicted alpha scattering	Only small deflections	Most pass through, few large deflections
Consistent with experiment?	No - cannot explain large-angle scattering	Yes

Discovery of Protons and Neutrons

Subsequent experiments led to the discovery of positive particles (protons) and neutral particles (neutrons) in the nucleus. Protons were discovered in 1919 by Rutherford and neutrons in 1932 by British scientist James Chadwick (1891-1972). Thus, the atom is composed of electrons, protons, and neutrons.

As chemists, we can take a simple view because only three subatomic particles—the proton, neutron, and electron—have a bearing on chemical behavior. The charge of an electron is -1.602 × 10^-19 C. That of a proton is equal in magnitude, +1.602 × 10^-19 C. The quantity 1.602 × 10^-19 C is called the elementary charge. Neutrons have no charge.

Subatomic Particle	Charge	Mass (g)	Mass (amu)	Location
Electron	1- ( -1.602e-19 C)	9.10938 × 10-28	0.00055	Outside nucleus
Proton	1+ (+1.602e-19 C)	1.67262 × 10-27	1.0073	In nucleus
Neutron	0	1.67493 × 10-27	1.0087	In nucleus

Callout Box:Size Matters > If an atom were the size of a football stadium, the nucleus would be a marble at the center, and the electrons would be like gnats buzzing around the stands. The nucleus is about 10,000 times smaller than the atom, yet contains 99.9% of its mass. This explains why most alpha particles miss it entirely.

Exam Trap:Students think Rutherford's experiment showed the nucleus is negative because alpha particles were repelled. Actually, alpha particles are positive. Repulsion means the nucleus is also positive. Like charges repel. Also, do not say atoms are mostly empty because electrons are tiny - it is mostly empty because the distance between nucleus and electrons is enormous relative to their sizes.

Practice Questions

MCQs (Multiple Choice Questions)

01. What did J.J. Thomson conclude about cathode rays?

1. They are electromagnetic waves
2. They are streams of negatively charged particles present in all matter
3. They are positively charged ions of the cathode material
4. They are neutrons

Correct: B
Explanation: Thomson observed that e/m ratio was independent of cathode material and gas, proving cathode rays are a universal negatively charged subatomic particle, later named electron. They are not waves, not positive, and neutrons were discovered later.

02. In Thomson's experiment, why were electric and magnetic fields applied perpendicular to each other?

1. To increase the speed of electrons
2. To balance forces so electrons travel straight and allow velocity measurement
3. To produce fluorescence
4. To ionize the gas

Correct: B
Explanation: Electric field deflects electrons one way, magnetic field the opposite. Balancing them gives straight-line motion, where eE = evB, allowing calculation of v = E/B, which then allows e/m calculation from deflection with one field alone.

03. What was the key observation in Millikan's oil-drop experiment that allowed determination of elementary charge?

1. All oil drops fell at same speed
2. Charge on drops was always integral multiple of 1.602 × 10^-19 C
3. Oil drops were positively charged
4. X-rays increased mass of drops

Correct: B
Explanation: Millikan found charges were quantized: q = n × e, where n is integer. The smallest common factor was e. This proved charge is quantized and gave its value.

04. Which radiation consists of electrons?

1. Alpha
2. Beta
3. Gamma
4. None

Correct: B
Explanation: Beta particles are high-speed electrons emitted from nucleus during neutron decay. Alpha are He nuclei, gamma are EM radiation with no mass or charge.

05. Why are beta rays deflected more than alpha rays in an electric field?

1. Beta has larger charge
2. Beta has much larger charge-to-mass ratio
3. Gamma interferes
4. Alpha is neutral

Correct: B
Explanation: Deflection a = F/m = qE/m. Beta q/m is ∼1.76×10⁸ C/g, alpha q/m is much smaller (∼4.8×10⁴ C/g). So for same field, beta accelerates more.

06. What did Rutherford observe in gold foil experiment that disproved plum-pudding model?

1. All alpha particles passed through
2. Some alpha particles were scattered at large angles >90°
3. Alpha particles were absorbed
4. Gold foil became radioactive

Correct: B
Explanation: Plum-pudding predicts only tiny deflections due to diffuse positive charge. Large-angle scattering requires concentrated positive charge - nucleus.

07. Most alpha particles pass through gold foil undeflected because:

1. Gold atoms are neutral
2. Atom is mostly empty space and nucleus is tiny
3. Alpha particles have no charge
4. Electrons absorb alpha particles

Correct: B
Explanation: Nucleus occupies ∼1/10,000 diameter of atom. Probability of direct hit is very low, so most alpha particles travel through empty space far from nucleus.

08. The mass of an electron is approximately:

1. Equal to proton mass
2. 1/2000 of hydrogen atom mass
3. 7400 times proton mass
4. Equal to neutron mass

Correct: B
Explanation: Electron mass 9.11×10^-28 g, hydrogen atom ∼1.67×10^-24 g, ratio ∼1/1836 ∼1/2000.

09. Who discovered the neutron and when?

1. Rutherford 1919
2. Chadwick 1932
3. Thomson 1897
4. Becquerel 1896

Correct: B
Explanation: James Chadwick discovered neutron in 1932 by bombarding beryllium with alpha particles. Rutherford discovered proton in 1919.

10. In Rutherford's model, where is the positive charge located?

1. Spread throughout atom
2. In electrons
3. In small dense central nucleus
4. In empty space

Correct: C
Explanation: Nuclear model places all positive charge and most mass in tiny nucleus, explaining strong repulsion at close approach.

FRQs (Free Response Questions)

FRQ 1: (AP Style)
Explain how Thomson's cathode-ray experiment determined that electrons are fundamental particles common to all matter. Include role of charge-to-mass ratio.

Rubric:

• Description of cathode-ray tube producing beam from cathode to anode, fluorescing (1 pt)
• Deflection by electric/magnetic fields shows negative charge, and balancing fields gives velocity (1 pt)
• 1 pt: e/m = 1.76×10⁸ C/g independent of cathode material/gas, proving universal particle (1 pt)
• Conclusion that because same ratio for all materials, electrons are subatomic component of all atoms (1 pt)

FRQ 2: (IB HL Style)
Describe Millikan's oil-drop experiment procedure and show calculation of electron mass using Thomson's ratio.

Rubric:

• Oil drops sprayed, fall through hole, ionized by X-rays to become charged (1 pt)
• Voltage varied to suspend drop, balancing mg = qE, measuring q (1 pt)
• Observation that q = n × 1.602×10^-19 C, quantizing charge (1 pt)
• Formula m = e / (e/m) correctly set up (1 pt)
• Calculation: 1.602e-19 / 1.76e8 = 9.10e-28 g, agreement with modern value (1 pt)

FRQ 3: (A-Level Style)
A beam containing alpha, beta and gamma radiation enters a uniform electric field between two charged plates. Sketch and explain the paths.

Rubric:

• Diagram: beta to positive plate large deflection, alpha to negative small deflection, gamma straight (1 pt)
• Explanation of charge: beta negative to positive, alpha positive to negative, gamma neutral (1 pt)
• Explanation of deflection magnitude related to q/m: beta large q/m → larger deflection (1 pt)
• Mention that alpha mass ∼7400× electron, so less deflection (1 pt)

FRQ 4: (College Chemistry 1)
Rutherford's gold foil experiment: 1 in 20,000 alpha particles scattered >90°. Explain how this led to nuclear model and estimate nucleus size.

Rubric:

• Most pass through → atom mostly empty (1 pt)
• Few large-angle → concentrated positive charge needed for strong repulsion, diffuse model fails (1 pt)
• Nuclear model: tiny dense nucleus with all positive charge and mass, electrons in empty space (1 pt)
• Implication that nucleus diameter ∼10^-14 m vs atom 10^-10 m, factor 10,000 (1 pt)

FRQ 5: (USNCO/IChO Style)
Compare Thomson plum-pudding and Rutherford nuclear models in terms of charge distribution, mass distribution, and predicted alpha scattering. Why did Rutherford need Marsden to look at large angles?

Rubric:

• Charge: Thomson uniform sphere, Rutherford concentrated nucleus (1 pt)
• Mass: Thomson uniform, Rutherford almost all in nucleus (1 pt)
• Scattering prediction: Thomson only small ∼1°, Rutherford predicts mostly none plus rare large >90° due to Coulomb repulsion (1 pt)
• Reason for large-angle search: completeness, because if only small angles looked, both models could appear consistent; large-angle is decisive test (1 pt)
• Concludes observation of back-scattering definitively disproves plum-pudding, proves nucleus (1 pt)

Frequently Asked Questions

What are cathode rays and why do they travel from cathode to anode?

Cathode rays are streams of electrons emitted from the negative cathode when high voltage is applied in a vacuum tube. They travel to the positive anode because opposite charges attract, and their path is visible via fluorescence on a screen inside the tube.

How did J.J. Thomson measure the charge-to-mass ratio of the electron?

Thomson balanced perpendicular electric and magnetic fields so the electron beam traveled straight. From field strengths he found velocity v=E/B, then measured deflection with only one field to get e/m=1.76×10⁸ C/g, proving electrons are subatomic.

What did Millikan's oil-drop experiment prove?

Millikan measured the elementary charge as 1.602×10^-19 C by suspending charged oil drops between plates. He showed charge is quantized in integral multiples of this value and used it with e/m to calculate electron mass 9.10×10^-28 g.

What are alpha, beta, and gamma radiation?

Alpha is a helium nucleus with 2+ charge and mass 7400× electron, beta is a high-speed electron with 1- charge, gamma is high-energy electromagnetic radiation with no charge or mass. Beta is deflected most in electric fields.

What is the plum-pudding model of the atom?

Proposed by Thomson, it described the atom as a uniform sphere of positive charge with electrons embedded like raisins in pudding. It predicted only small alpha deflections and was disproved by Rutherford's large-angle scattering.

What happened in Rutherford's gold foil experiment?

Most alpha particles passed straight through gold foil, some deflected slightly, and about 1 in 20,000 bounced back at large angles. This showed atoms are mostly empty space with a tiny dense positive nucleus that repels alpha particles.

Why is Rutherford's model called the nuclear model?

Because it places all positive charge and almost all mass in a central nucleus that is ∼10,000 times smaller than the atom, with electrons moving in the empty space around it. This explains why most alpha particles miss the nucleus.

What is the difference between proton, neutron and electron in location and charge?

Protons (+1) and neutrons (0) are in the nucleus with mass ∼1 amu. Electrons (-1) orbit outside the nucleus with mass ∼1/2000 amu. Proton and electron charges are equal magnitude 1.602×10^-19 C.

Why do particles with same charge repel and opposite charges attract?

This is a fundamental electrostatic law observed in cathode ray and radioactive deflection experiments. It underlies atomic structure: positive nucleus attracts negative electrons but repels positive alpha particles, causing scattering.

Exam Mapping

Concept	A-Level Chemistry (AQA/OCR/Edexcel)	AP Chemistry	IB Chemistry HL	College Chemistry 1	USNCO / IChO
Cathode rays & electron discovery	3.1.1.1 Atomic structure - history of atom, electron discovery	Big Idea 1 - Structure of atom, electron	2.1 The nuclear atom - evidence for electrons	Ch 2: Thomson experiment, e/m	History of atomic theory, interpretation of cathode ray data
Millikan oil-drop & elementary charge	Calculation of e/m and charge quantization	Laboratory question - charge quantization	Nature of science - Millikan	Quantitative - calculating mass from e/m	Quantization, error analysis
Radioactivity α/β/γ properties	3.1.1.1 Types of radiation, penetration, deflection	Nuclear chemistry - types, balancing	11.2 Nuclear reactions - properties	Ch 2: Radioactivity as evidence	Nuclear decay, deflection in fields, mass/charge
Plum-pudding vs nuclear model	Comparison, limitations of models	Models of atom evolution	2.2 Electron configuration - model development	Evolution of atomic model	Model evaluation, experimental design
Rutherford gold foil experiment	Core practical - scattering, nuclear atom	Experimental evidence for nucleus, empty space	2.1 - Rutherford experiment conclusions	Ch 2: Nuclear model, size of nucleus	Large-angle scattering, Coulomb's law, nucleus size estimate
Protons, neutrons, electrons charges/masses	Subatomic particle table, elementary charge	Atomic structure fundamentals	2.1 Relative masses/charges	Ch 2: Structure of atom	Memorization + calculations

Conclusion

The discovery of atomic structure transformed the atom from an indivisible billiard ball into a complex, mostly empty entity with a dense heart. Thomson's cathode-ray tube revealed the first subatomic particle, the electron, with its enormous charge-to-mass ratio. Millikan's oil drops quantized charge and gave us mass. Becquerel and the Curies uncovered radioactivity, providing both new particles and the perfect projectiles for Rutherford. Rutherford's gold foil experiment then demolished the plum-pudding model: most alpha particles sailing through empty space, but a few repelled violently, proved that all positive charge and almost all mass live in a nucleus ten thousand times smaller than the atom itself.

Mastering these four experiments - how they were set up, what was observed, and why the observation forced a new model - gives you the framework for every atomic structure question on AP, A-Level, IB HL, College Chemistry, USNCO and IChO. From here, the next step is counting protons, neutrons, and electrons and understanding isotopes and atomic masses.

References

Brown, LeMay, Bursten, Murphy, Woodward, Stoltzfus - Chemistry: The Central Science, 15th Edition, Chapter 2.2 The Discovery of Atomic Structure: Cathode Rays and Electrons, Radioactivity, The Nuclear Model of the Atom. Pearson.