The History of The Atomic Model
The atom is the smallest unit of matter that makes up everything around us. Different types of atoms form the elements of the periodic table. While each element has its own unique properties, every atom contains the same subatomic particles: protons, neutrons, and electrons.
However, it took hundreds of years of careful observation, experimentation, and revision for scientists to figure out what atoms are made of and how they are arranged. Today, our most accurate description is called the quantum mechanical model of the atom.
Let’s explore the history of the atomic model and some of the scientists who helped shape our current understanding.
Democritus, 450 BC
Long before scientific labs existed, philosophers tried to explain the nature of matter.
Democritus, a Greek philosopher, argued that matter could not be divided forever — there must be a smallest piece. He called this smallest piece of matter the atomos, meaning "indivisible." A generation later, Aristotle rejected the idea. He believed that you could divide a grain of sand (or any matter) infinitely. Aristotle's view won out and dominated Western thought for the next two thousand years, which is part of why Democritus's idea had to wait so long to be taken seriously again.
An infographic showing the history of changes to the atomic model. From solid spheres to electron clouds: five models that trace how our picture of the atom evolved as each new discovery forced scientists to revise what came before.
Atomic theory can feel like a list of names and dates to memorize. This graphic organizer turns it into a story, one where each scientist builds on, corrects, or refines the discovery before them, and where the model of the atom slowly comes into focus.
Students work through ten key figures, from Dalton's solid spheres to the quantum model, recording what each scientist discovered, the evidence or experiment behind it, and how the picture of the atom changed as a result. Some rows are marked "No New Model" on purpose — a lesson that not every discovery redraws the atom, and a natural way to reinforce the Nature of Science.
The organizer pairs with my free blog post "The History of the Atomic Model," included as a QR code on every page so students can pull up the reading on any device.
What's Included:
✔ Complete Version — the filled-in answer key, ready to use as a reference, a modeling tool, or a self-check station
✔ Blank Version — scientist names provided as anchors; students complete the rest from the reading
✔ Cut & Paste Version — years and select cells pre-filled, with ready-to-cut card sets for discoveries, evidence, and atomic models
✔ Card sets: 10 discovery cards, 10 evidence cards, and 5 labeled model cards
✔ Instructor Notes with differentiation guidance and printing tips
✔ QR code linking directly to the companion reading
✔ Terms of Use and Credits pages
Why teachers like it:
Three versions means three levels from one activity. Assign the blank organizer to on-level students, hand the cut & paste version to students who need scaffolding, and keep the answer key for you, your co-teacher, or your sub folder.
Great for:
Introducing atomic structure or atomic theory
Sub plans and station rotations
Review before an atomic structure test
Interactive notebooks
Homework or flipped-classroom reading
Want to try it first? Grab the free blank organizer from my store, then come back for the answer key and cut & paste version.
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John Dalton, 1804
John Dalton, an English schoolteacher and chemist, revived the idea of atoms with one of the first true atomic theories. By studying the way elements combined in fixed ratios (the law of definite proportions or Proust’s Law) and connecting it with the law of conservation of mass, he proposed:
All matter is made of atoms.
Atoms of the same element are identical.
Atoms cannot be created, divided, or destroyed.
Atoms of different elements can combine in simple ratios to form compounds.
Atoms of the same element can combine in more than one ratio to form two or more compounds.
In chemical reactions, atoms are rearranged.
In addition to his postulates, Dalton pictured atoms as solid spheres, like tiny billiard balls, with no internal structure. This became known as the solid sphere model.
We now know that parts of Dalton’s theory are not correct. Atoms can be subdivided (into protons, neutrons, and electrons), isotopes show that atoms of the same element can have different masses, and nuclear reactions can change atoms into other elements. But Dalton’s work was a crucial first step in building a scientific model of the atom.
Michael Faraday, 1830s
Michael Faraday, an English scientist, studied the relationship between electricity and matter. Through his experiments on electrolysis (using electricity to break down compounds), Faraday discovered that atoms and electricity were connected. He showed that electrical forces play a role in holding compounds together and even coined words we still use today: electrode, cathode, anode, and ion.
Faraday did not discover subatomic particles, but his work suggested that atoms themselves must contain electrical components. This insight laid the groundwork for experiments with cathode rays later in the 1800s, which eventually revealed the electron.
Cathode Ray Tubes
After Faraday’s work suggested that electricity was connected to atoms, scientists began experimenting with devices called cathode ray tubes. A cathode ray tube is a sealed glass tube with almost all of the air removed. When electricity is passed through it, a glowing ray travels from the cathode (negative end) to the anode (positive end).
At first, no one knew what this mysterious “ray” was. But experiments showed that the beam could cast shadows, move small paddles, and bend when magnets or electric fields were nearby. This meant that the rays had both mass and charge.
These cathode rays would become the focus of J. J. Thomson’s experiments.
Image Credit: Figure 2.6. Thomson’s cathode-ray experiments. Adapted from Chemistry 2e: 2.2 Evolution of atomic theory (OpenStax, 2019), https://openstax.org/books/chemistry-2e/pages/2-2-evolution-of-atomic-theory, CC BY 4.0. Image credits: modification of work by Nobel Foundation; modification of work by Eugen Nesper; modification of work by Kurzon/Wikimedia Commons.
J. J. Thomson, 1897
By the 1890s, cathode rays had been studied for decades, however there were still many questions surrounding it and the ray. J. J. Thomson, working at the Cavendish Laboratory in Cambridge, ended the argument by measuring.
First, Thomson showed the rays carry a negative charge. Thomson removed most of the gas in the chamber, and then placed to oppositely charged electric plates around the cathode ray. The beam bent away from the negatively charged plate, and pulling toward the positive plate. Whatever the rays were made of, it was negatively charged.
Second, he measured how heavy the particles were. By balancing an electric field against a magnetic field, Thomson could calculate the ratio of a particle's charge to its mass. The result was startling: these particles were more than a thousand times lighter than a hydrogen ion, the lightest thing anyone had ever weighed. (We now know the electron is about 1,836 times lighter than a proton.)
Third, he showed the results were universal. Thomson swapped out the electrode materials. He then changed the gas left in the tube. The charge-to-mass ratio came out the same every time. If the same tiny particle comes out of every material you test, it isn't a property of aluminum or platinum. Atoms have parts, and every atom has the same parts.
Dalton's claims of an indivisible atom no longer held true, ninety years after he proposed it. Thomson had identified the first subatomic particle. This particle was later named the electron.
That left Thomson with a new problem. Electrons are negatively charged, but atoms are electrically neutral overall. So there had to be positive charge somewhere in the atom to balance them out, and something holding the whole thing together. His answer was the plum pudding model: a sphere of positive charge with the tiny electrons scattered throughout it, like plums in a pudding. Thomson received the Nobel Prize in Physics in 1906 for this work. His model didn't survive the decade, but the electron did and it was the first idea to show that atoms were not indivisible.
Image Credit: Figure 2.8. Early atomic models: (a) Thomson’s “plum pudding” model; (b) Nagaoka’s Saturn-like model. Adapted from Chemistry 2e: 2.2 Evolution of atomic theory (OpenStax, 2019), https://openstax.org/books/chemistry-2e/pages/2-2-evolution-of-atomic-theory, CC BY 4.0. Image credits: modification of work by Man vyi/Wikimedia Commons; modification of work by NASA/Wikimedia Commons.
Robert Millikan, 1909
Robert Millikan, an American physicist, measured the charge of the electron with his famous oil drop experiment. By balancing the downward pull of gravity with the upward electric force on tiny charged oil droplets, Millikan determined the charge of an electron. This also allowed scientists to calculate its mass.
Image Credit: Figure 2.7. Millikan’s oil-drop experiment. Adapted from Chemistry 2e: 2.2 Evolution of atomic theory (OpenStax, 2019), https://openstax.org/books/chemistry-2e/pages/2-2-evolution-of-atomic-theory, CC BY 4.0.
Ernest Rutherford, 1911
Ernest Rutherford, often called the "father of nuclear physics," directed the gold foil experiment, carried out in his lab by Hans Geiger and Ernest Marsden. He shot positively charged alpha particles at a thin sheet of gold foil. Most particles passed straight through, but a few bounced back at sharp angles. This was unexpected — if mass and positive charge were spread evenly throughout the atom, as the plum pudding model suggested, the alpha particles should have passed through with only tiny deflections. None should have bounced back. Rutherford famously described the result as being "as if you had fired a 15-inch [artillery] shell at a piece of tissue paper and it came back and hit you."
This surprising result showed that atoms are mostly empty space, with a small, dense, positively charged center called the nucleus. Rutherford's model placed the electrons somewhere outside the nucleus, though he said little about how they were arranged. This new model of the atom became known as the Nuclear Model.
Image Credit: Figure 2.9. Rutherford’s gold-foil experiment. Adapted from Chemistry 2e: 2.2 Evolution of atomic theory (OpenStax, 2019), https://openstax.org/books/chemistry-2e/pages/2-2-evolution-of-atomic-theory, CC BY 4.0.
Henry Moseley, 1913
Henry Moseley, an English physicist, studied X-rays emitted by different elements. He discovered that each element has a unique amount of positive charge in its nucleus, now known as the atomic number.
Moseley's work corrected the arrangement of the periodic table, proving that elements should be ordered by atomic number rather than atomic mass.
Niels Bohr, 1913
Rutherford's nuclear model had a flaw. According to the physics of the time, an electron circling a nucleus should continuously radiate energy, be unstable, and collapse. Atoms, however, don't do this.
Bohr's solution was to change the rules rather than the picture. Borrowing from Max Planck and Albert Einstein — who had shown that energy comes in discrete packets called photons rather than any amount you like — Bohr proposed that an electron can only occupy certain allowed orbits, each with a fixed energy. Electrons can still move between levels, absorbing energy to jump up and releasing it as light when they fall back down.
This became known as the Bohr Model of the atom. Because its neat orbits look like a miniature solar system, Bohr's model is often nicknamed the planetary model (though you'll sometimes see that nickname applied to Rutherford's model too, since Bohr kept Rutherford's nucleus and added the energy levels).
Bohr’s Model successfully explained the spectral lines of hydrogen, or the colors emitted by hydrogen. However, this model didn’t work for other elements. Because of this limitation, Bohr’s model was later replaced by more advanced models to better explain the structure of atoms.
However, we still often use Bohr’s model to represent atoms today, especially in introductory chemistry classes! Models can still be useful representations even though they may not be accurate.
Ernest Rutherford (Again!), 1917 - 1920
Rutherford wasn't finished after discovering the nucleus. In a new series of experiments, he fired alpha particles at nitrogen gas and hydrogen nuclei came flying out. Since a hydrogen nucleus is the simplest one there is (commonly just a single positive particle, with no neutrons), this could only mean one thing: hydrogen nuclei must exist inside the atoms of other elements. They were a fundamental building block of matter.
Rutherford named this particle the proton, from the Greek protos, meaning "first." At last, scientists knew what gave the nucleus its positive charge and Moseley's atomic number gained a physical meaning: it counts the protons in an atom's nucleus. As a bonus, by having a proton be emitted from the nitrogen, Rutherford had accidentally transformed it into oxygen — the first time humans ever changed one element into another.
Louis de Broglie and Erwin Schrödinger, 1920s
Louis de Broglie proposed in 1924 that electrons have a dual nature — they behave as both particles and waves. This had an immediate payoff: a wave confined to a loop can only exist at certain wavelengths. The wave has to meet up with itself after going around, meaning that only certain orbits are possible. Bohr's energy levels stopped being an assumption and became a consequence.
In 1926, Erwin Schrödinger turned that idea into mathematics, and Max Born showed how to read the solutions. These solutions were not the electron's location, but the probability of finding it in a given place. In Bohr's model, the electron is a ball on a track. In this model, there is no track. Electrons are constantly moving, and they do not have a fixed or defined position within the atom.
Because of that, what we draw instead is an orbital — the region where an electron will most probably be found. The s, p, d, and f orbital shapes you learn in chemistry class are solutions to Schrödinger's equation. Every time you write an electron configuration, you're using quantum mechanics.
This became the quantum mechanical model of the atom. In this model, electrons are not pictured as moving in fixed orbits, but as existing in regions of probability.
📦 Fun Fact: Schrödinger’s Cat
In 1935, Schrödinger created a thought experiment where he described a cat sealed in a box with a poison gas triggered by a decaying radioactive atom. Until someone opens the box, the standard interpretation of quantum mechanics would say the cat is both alive and dead.
Here's the part that usually gets left out: Schrödinger wasn't presenting this as a wonder of physics. He was critiquing shortcomings that he saw with the current interpretations of quantum physics. His point was that if you take the prevailing interpretation seriously and scale it up to something the size of a cat, you get a conclusion nobody can accept, so something must be wrong with the interpretation. The most famous thought experiment in physics was written as a complaint.
While this thought experiment wasn’t about atoms directly, it highlighted the same quantum ideas Schrödinger used in his quantum model.
James Chadwick, 1932
Up until this point, scientists had discovered two subatomic particles, the proton and the electron. However, it had been speculated that a neutral particle may exist. But, it wasn’t confirmed until James Chadwick discovered the neutron, an uncharged particle located in the nucleus alongside protons. Neutrons explained why atoms of the same element can have different masses, also known as isotopes.
Chadwick read the experimental accounts of other scientists searching for the neutron and conducted experiments of his own. He bombarded beryllium with alpha particles and found it gave off a strange, highly penetrating radiation that carried no charge. Others had assumed it was gamma rays. Chadwick aimed it at paraffin wax and measured the protons it knocked loose, and the energies were far too high for massless gamma rays to produce. Only a neutral particle with roughly the mass of a proton could transfer that much. He called it the neutron.
With the discovery of neutrons, the modern picture of the atom — protons and neutrons in the nucleus with electrons in surrounding orbitals — was in place.
Look at the path: a solid sphere, a pudding with electrons in it, a nucleus with electrons around it, fixed energy levels, and finally probability clouds. Throughout the years, there have been many iterations of the atomic model, and every single time it changed, it was because someone found evidence the old model couldn't explain.
What makes the quantum model different is that a century later, nobody has discovered a more accurate model. We've refined it and filled in details, but the model still holds. If new evidence is presented, then scientists will continue to use the evidence and prior research to develop our understanding of the atom.
