Historical Development of Atomic Models

Why Atomic Models Were Needed

As chemistry and physics developed, scientists realized that matter must have an internal structure that explains:

• Why substances combine in fixed proportions
• Why elements emit characteristic colors of light
• Why some “rays” pass through matter and others are strongly absorbed
• Why atoms are neutral but can form charged particles (ions)

Atomic models are simplified pictures that try to explain such observations. Over time, these models were revised as new experiments revealed contradictions with older ideas.

This chapter traces how views of the atom evolved up to the point where quantum ideas became necessary, preparing for later discussion of the Bohr–Sommerfeld and modern quantum mechanical models.

Early Ideas of Atoms (Antiquity to 18th Century)

Long before modern science, philosophers speculated about the divisibility of matter.

• Greek atomism (Democritus, ~5th century BCE)
• Idea: Matter is made of tiny, indivisible particles called “atoms” (from “atomos” = uncuttable) moving in empty space.
• Atoms differ in shape and size; macroscopic properties arise from arrangements of atoms.
• No experiments supported this; it was a philosophical speculation, not a scientific theory.
• Continuum view (e.g., Aristotle)
• Idea: Matter is continuous and composed of a few “elements” (earth, water, air, fire).
• Dominated Western thought for centuries and shaped alchemical traditions.

These early notions set the stage but did not yet form a quantitative scientific model of atoms.

Dalton’s Atomic Theory (Early 19th Century)

By the late 18th and early 19th centuries, careful measurements in chemistry (especially mass measurements) showed regularities in how substances combine.

Experimental Foundations

Key laws that emerged from chemical experiments:

• Law of conservation of mass (Lavoisier, late 18th century)
  In a closed system, total mass remains constant during a chemical reaction.
• Law of definite proportions (Proust)
  A given compound always contains the same elements in the same mass ratio, regardless of its source or preparation.
• Law of multiple proportions (Dalton)
  When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in simple whole-number ratios.

These patterns suggested that matter consists of discrete building blocks.

Dalton’s Model

John Dalton (1803–1808) proposed a scientific atomic theory to explain these laws:

• Postulates (simplified)
• Matter is made of tiny, indivisible particles called atoms.
• Atoms of a given element are identical in mass and properties; atoms of different elements differ.
• Compounds form when atoms of different elements combine in simple whole-number ratios.
• Chemical reactions are rearrangements of atoms; atoms are not created or destroyed in reactions.
• Model image
• Atoms are solid, indivisible spheres, like tiny billiard balls.

Successes and Limitations

• Successes
• Explained conservation of mass as conservation of atoms.
• Explained definite and multiple proportions using fixed numbers of atoms in compounds.
• Gave a rational basis for chemical formulas and stoichiometry (treated separately).
• Limitations
• Atoms were assumed indivisible and structureless.
• No internal components (such as electrons) were recognized.
• Could not explain electrical phenomena or spectroscopy.

Despite limitations, Dalton’s picture established atoms as real, countable entities in chemistry.

Evidence for Subatomic Particles: Discovery of the Electron

By the late 19th century, work with electricity and “cathode rays” showed that atoms were not indivisible.

Cathode Ray Experiments (J. J. Thomson, 1897)

• Setup
• A glass tube with low-pressure gas and two electrodes.
• High voltage applied → a beam (cathode rays) travels from the negative electrode (cathode) to the positive one (anode).
• The beam causes the glass to glow where it strikes.
• Key observations
• The rays are deflected by electric and magnetic fields, indicating they carry charge.
• Deflection direction showed the charge is negative.
• Measuring the deflection allowed determination of the charge-to-mass ratio, $e/m$.
• Conclusions
• Cathode rays are streams of negatively charged particles.
• These particles are the same regardless of the gas in the tube or electrode material.
• Therefore, they are universal components of atoms.

J. J. Thomson interpreted these particles as electrons, the first subatomic particles to be identified.

Millikan’s Oil Drop Experiment (1909–1911)

Robert Millikan measured the charge of a single electron by:

• Suspending tiny charged oil droplets between electrically charged plates.
• Adjusting the voltage until droplets were held stationary.
• Using the force balance to calculate the charge on individual droplets.

Key outcome:

• The observed charges were all integer multiples of a smallest charge value, interpreted as the elementary charge $e$ (charge of the electron).
• Combined with Thomson’s $e/m$ ratio, this gave the electron’s mass, much smaller than that of a hydrogen atom.

Together, these experiments showed:

• Atoms contain negatively charged, very light electrons.
• Atoms must also contain positive charge to be electrically neutral.

This contradicted Dalton’s idea of indivisible, homogeneous atoms.

Thomson’s “Plum Pudding” Model

To accommodate electrons while keeping overall neutrality, J. J. Thomson proposed a new atomic model around 1904.

Basic Picture

• The atom is a diffuse sphere of positive charge.
• Electrons are embedded in this sphere like “plums in a pudding” or raisins in a cake.
• The total negative charge of electrons balances the total positive charge.

Properties and Aim

• Intended to explain:
• Electrical neutrality of atoms.
• The existence of negatively charged electrons inside atoms.
• The positive charge was spread out, not concentrated in a central nucleus.

Shortcomings

Although plausible at first, this model would soon fail when radioactive phenomena and scattering experiments revealed concentrated positive charge in atoms.

Radioactivity and the Path Toward a Nuclear Model

Around the turn of the 20th century, new discoveries complicated the picture of atomic structure.

Discovery of Radioactivity (Becquerel and the Curies, 1896–1900)

• Certain elements (e.g., uranium, radium) spontaneously emit penetrating rays.
• These emissions are of different types, later classified as:
• Alpha ($\alpha$) radiation: positively charged particles.
• Beta ($\beta$) radiation: negatively charged particles (electrons).
• Gamma ($\gamma$) radiation: high-energy electromagnetic radiation.

Radioactivity showed that atoms can transform into other elements and emit subatomic particles, challenging the notion of permanent, indivisible atoms.

Scattering Experiments as a Probe of Atomic Structure

Ernest Rutherford and collaborators used $\alpha$ particles (helium nuclei) as projectiles to probe how matter is organized inside atoms. These experiments were decisive in overturning the plum pudding model.

Rutherford’s Nuclear Model (1911)

The gold foil experiment (Geiger, Marsden, and Rutherford) provided direct evidence for a concentrated atomic center.

The Gold Foil Experiment

• Setup
• A thin gold foil bombarded by a beam of $\alpha$ particles.
• A detection screen around the foil observed where $\alpha$ particles struck.
• Expectations from the plum pudding model
• $\alpha$ particles, being massive and positively charged, should pass through with only very small deflections.
• The positive charge in atoms was thought to be spread out, causing only gentle scattering.
• Observations
• Most $\alpha$ particles passed through the foil with little or no deflection.
• A small fraction were deflected at large angles.
• A very small number were deflected backward, nearly 180°.

Rutherford famously compared the large-angle scattering to firing a shell at tissue paper and having it bounce back.

Interpretation

To explain such strong deflections:

• The positive charge could not be spread out.
• Instead, it had to be concentrated in a very small, massive region.

Rutherford proposed the nuclear model:

• The atom consists of a tiny, dense, positively charged nucleus.
• Most of the atom is empty space.
• Negatively charged electrons move around the nucleus.

Features and Consequences

• Positive charge and most of the mass are in the nucleus.
• Electron cloud occupies most of the volume but contributes little mass.
• The model explained:
• Strong deflections of $\alpha$ particles due to close encounters with a concentrated positive center.
• Why most $\alpha$ particles pass straight through (they miss the tiny nucleus).

However, the nuclear model still left key questions:

• How are electrons arranged around the nucleus?
• Why are atoms stable if orbiting electrons should radiate energy?
• How to explain discrete atomic spectra?

These questions led to further refinement.

Development of the Nuclear Concept: Protons and Neutrons

The idea of the nucleus quickly evolved as new experiments clarified its composition.

Identification of the Proton

• Hydrogen nucleus as simplest case
• Hydrogen has atomic number 1 and is the lightest element.
• Rutherford’s further work on scattering and nuclear reactions indicated that:
• The nucleus of hydrogen is a single, positively charged particle.
• This particle, with charge $+e$ and relatively large mass (compared to the electron), became known as the proton.
• The charge of a nucleus with atomic number $Z$ is then $+Ze$.

The Mass Problem and the Neutron Hypothesis

It was observed:

• Atomic mass is roughly proportional to, but generally larger than, the atomic number $Z$.
• Early models assumed nuclei contained protons and electrons, but:
• This was inconsistent with observed nuclear properties and quantum theory.
• It raised difficulties explaining nuclear spin and stability.

A neutral nuclear particle was hypothesized to resolve these issues.

Discovery of the Neutron (James Chadwick, 1932)

• Experiments
• Bombarding beryllium and other light elements with $\alpha$ particles produced highly penetrating radiation.
• This radiation was not deflected by electric or magnetic fields, indicating no net charge.
• However, it could knock protons out of other materials, showing it had mass and could transfer momentum.
• Conclusion
• The radiation consisted of neutral particles with mass similar to the proton.
• These particles were named neutrons.

Nuclear Composition in Modern Terms

After these developments, the nucleus was understood as:

• Composed of protons (positively charged) and neutrons (neutral).
• Together, protons and neutrons are called nucleons (treated in detail elsewhere).

This clarified:

• The atomic number $Z$ as the number of protons.
• The mass number $A$ as the total number of protons and neutrons.

From Nuclear Model to Quantum Models

While Rutherford’s nuclear model correctly described the atom’s central structure, it did not explain:

• Why electrons do not spiral into the nucleus due to radiative energy loss.
• Why atoms emit and absorb light only at discrete wavelengths (line spectra).

These problems inspired:

• Bohr’s model (1913), introducing quantized electron orbits to explain the hydrogen spectrum.
• Sommerfeld’s refinements, and eventually
• The modern quantum mechanical model with orbitals and probability distributions.

These later models provided a more complete description of atomic structure and will be developed in subsequent chapters. Here, it is important to see them as part of a historical progression:

1. Philosophical atoms →
2. Dalton’s indivisible spheres →
3. Thomson’s electron-containing “plum pudding” →
4. Rutherford’s nuclear atom →
5. Proton–neutron nucleus with quantum-mechanical electrons.

Summary of the Historical Development

• Early atomism introduced the idea of indivisible particles but lacked experimental grounding.
• Quantitative chemistry in the 18th–19th centuries led to Dalton’s atomic theory, which explained fixed composition laws.
• Electrons were discovered via cathode ray experiments, proving atoms contained smaller, charged constituents.
• Thomson’s plum pudding model attempted to incorporate electrons into an overall neutral atom.
• Radioactivity and scattering experiments, especially Rutherford’s gold foil experiment, revealed a small, dense, positively charged nucleus.
• Identification of protons and discovery of neutrons completed the basic picture of nuclear composition.
• The inability of classical nuclear models to explain atomic stability and spectra led to the development of quantum-based atomic models, which will be treated in later chapters.