The Basic Structure of the Atom

Chemistry and Our Universe

In this lecture, you will learn how each of these particles was discovered and how the culmination of generations of scientific work ultimately came together to create our understanding of the most fundamental unit of matter: the atom.

Atoms are comprised of 3 types of particles called subatomic particles: positively charged protons, negatively charged electrons, and neutrons, which have no charge. Protons and neutrons are of nearly equal mass and reside in a dense nucleus at the center of the atom. But much smaller, negatively charged electrons orbit this nucleus, balancing out the positive charge provided by the protons.


  • The notion of atoms was first forwarded by ancient Greek philosophers who postulated that there are just a handful of fundamental substances that combine in various ways to form all others. The Greeks widely believed these elements to be air, water, earth, and fire—and they also believed that particles of these elements were absolutely indivisible.
  • Today, we know that none of these are actually elements, and we also know that atoms of true elements can in fact be divided, but that they change to a new element when they are. Nonetheless, the Greek term “atom,” meaning “not divisible,” has stuck and is still how we refer to the smallest quantity of a given element that can exist.


  • Discourse over the nature of matter and its fundamental units slowed to nearly a halt with the fall of ancient Greece, and it wasn’t until the start of the 19th century that the concept of the atom was revived, primarily as a result of the work of John Dalton. Dalton didn’t have the sophisticated scientific instrumentation that we do today, so he had no way to see or experiment directly with atoms. What he did have, though, was a keen intellect and the benefit of the work of Antoine Lavoisier and others.
  • Using relatively simple observations, Dalton was able to formulate a sound atomic theory. He found that when 2 elements combine to form a compound, the mass of all the products is equal to the masses of all the starting materials. This is the law of conservation of mass.
  • He also found that when 2 elements combine to form more than one compound, the weights of one of the elements that combine with a fixed weight of the other are in a ratio of very simple whole numbers. This is the law of multiple proportions.
  • It’s the combination of these 2 laws that gave Dalton an airtight argument for the existence of indivisible atoms that come together in these simple whole-number ratios to create molecules.
  • But Dalton had no way to effectively probe the structures of his proposed atoms. His observations proved their existence, but they could not explain exactly what atoms themselves were made of. We would have to wait another century before technology caught up with the atom, giving us a way to learn the inner workings of this remarkable construct of nature.


  • The first of these advances is the cathode-ray tube, which was developed at the end of the 19th century and was used in televisions up until a decade ago. Cathode-ray tubes played an important role in the advancement of human understanding of atoms.
  • Using the cathode-ray tube, J. J. Thomson was able to create isolated beams of pure electrons, measuring their mass, velocity, and charge. Thomson’s measurements led him to a startling conclusion: that electrons were much smaller than atoms.
  • This led Thomson to propose the first structural model for an atom. More than 100 years after Dalton repopularized the concept of atoms, the first attempt to explain their structure had finally been offered. Thomson postulated that atoms consisted of the small, negatively charged electrons he had observed—embedded in a very low-density, positively charged, spherical matrix making up the rest of its mass.
  • A sketch of this model, with electrons peppered onto a positively charged sphere, evokes images of raisins in a bowl of pudding, earning it the moniker of Thomson’s plum pudding model.
  • But this first attempt to explain atomic structure was short-lived, because less than 2 decades after Thomson proposed the plum pudding model, his protégé, Ernest Rutherford, conducted an experiment that intended to build on Thomson’s model, but instead disproved it.


  • Rutherford is credited with the discovery of what are known as alpha particles, which are much larger charged particles (more than 1000 times larger than electrons) created during the radioactive decay of uranium. But Rutherford is most famous not for discovering these particles, but for what he did with them next.
  • Rutherford was curious to know how these relatively large, energetic particles would interact with atoms as they passed through a thin foil, so he pointed a beam of alpha particles at a piece of gold foil, encircling the foil in a special fluorescent screen that would light up when struck with an alpha particle. Assuming that Thomson’s model was correct, the alpha particles were expected to simply pass through the foil, being scattered only slightly, if at all.
  • As the experiment progressed, he noted that the brightest spot was directly behind the gold foil, just as he had expected. But Rutherford was stunned when he turned his attention to the other side of the screen and saw that a few particles were deflected almost directly back toward the source.
  • Clearly, Thomson’s model was not correct. If the mass of gold atoms was distributed across their entire volume, then there is no way that an alpha particle could be reflected back toward the source.
  • Rutherford’s results strongly indicated that atoms were made of mostly empty space, with a highly concentrated nucleus containing most of its mass in just a small fraction of the atom’s total volume. This would explain why, on rare occasion, an alpha particle bounced back. Only an extremely dense point of matter taking up a small volume of the atom’s total could possibly withstand the impact of an alpha particle and cause it to ricochet back in its original direction.
  • So, there had to be an extraordinarily dense, but vanishingly small, point of mass at the center of atoms. Rutherford realized, then, that atoms were mostly made of empty space with a dense point of mass at their center. His model that accounts for this has been dubbed the Rutherford model, sometimes also called the nuclear model, because it is the first to acknowledge that most of the atom’s mass resides in a small, dense nucleus at its center.
  • Rutherford had discovered that atoms consisted of a dense, positively charged nucleus surrounded by very light, negatively charged electrons. While atoms have diameters of around 100 picometers, or about one-tenth of a nanometer, Rutherford’s work eventually led to the discovery that the radius of a typical nucleus is only about 1/100,000th that of its electron cloud—that is, 1 femtometer, or 0.000001 nanometers.
  • Rutherford had determined that massive, positively charged particles were concentrated in the nucleus of the atom. Protons had been discovered.


  • As research on radioactivity continued, there were some observations about atomic nuclei that weren’t adding up. Specifically, researchers had irradiated beryllium with alpha particles to produce a new kind of radiation—one that didn’t bounce off of nuclei like Rutherford’s alpha particles, but instead crashed into them with force, knocking protons out of target atoms.
  • The final key to understanding these unusual observations came in 1932, when James Chadwick, formerly a student of one of Rutherford’s assistants, completed the inventory of subatomic particles. Chadwick reasoned that the radiation from beryllium was able to penetrate other atomic nuclei better because it was uncharged.
  • Whereas the alpha particles used by Rutherford carried positive charge and were repelled by the positive nuclei of atoms they struck, sending them back, these new particles, though massive, must be uncharged.
  • The lack of an electrostatic repulsion allowed these new particles to strike other nuclei with great force—a force that Chadwick used to estimate the mass of one of these new particles, which he called neutrons. Not surprisingly, they were just a fraction of a percent larger than protons.
  • Because protons and neutrons are so similar in mass and contribute nearly all of the mass of an atom, chemists use a unit of mass called the atomic mass unit (amu) to compare the masses of atoms to one another.


  • By 1932, the combined thought of the ancient Greeks, Dalton, Thomson, Rutherford, Chadwick, and many others had led us to an understanding of an atomic structure that looks something like what we see today in popular depictions of the atom.
  • All atoms consist of a dense central core of matter, called the nucleus. In that nucleus are one or more protons. In the case of hydrogen, there is just one. An atom of the element hydrogen is the simplest atom.
  • Around that dense, positively charged core are electrons. In the case of hydrogen, there is just one electron to balance the positive charge from the proton in the nucleus. This is a complete hydrogen atom with a nucleus and an electron cloud around it. We sometimes refer to this version of hydrogen as protium, because its nucleus contains just a single proton.
  • The proton count in the nucleus is what gives elements their identity, so we can add a neutron to the hydrogen atom and still have hydrogen. But this makes the hydrogen atom about twice its original mass. We call this type of hydrogen—one in which a neutron is also present, making it more massive—deuterium.
  • We can increase the mass of the hydrogen atom again by adding another neutron to get tritium—another form of hydrogen that is 3 times the mass of its simplest form. All 3 of these atoms are hydrogen, but they are different by way of their masses. We call these atoms isotopes of one another.
  • If we add another proton to the atom’s nucleus and another electron to its cloud to balance out the charge, we now have a new element, because the number of protons in the nucleus has changed.
  • We can estimate the atomic mass of an elemental isotope simply by adding up the number of protons and neutrons in the particular isotope.
  • There are many elements that exist in nature as 2 or more isotopes. In fact, most of them do. When this is the case, the periodic table reports the average mass of a given element in nature.


  • When the number of protons and electrons are out of balance, we create a species that carries a net overall charge. These charged species can have properties that are drastically different than their corresponding neutral atoms.
  • These charged atoms are called ions. In general, if there are excess electrons, we get a negatively charged ion called an anion. If there is instead an excess of protons, we get a positively charged ion called a cation.



1 If negatively charged electrons are attracted to positively charged atomic nuclei, what keeps the electron cloud of an atom or ion from collapsing?

2 How many neutrons, protons, and electrons are in the following atoms and ions?

a An atom of plutonium 244

b An oxide ion (O2−) made from oxygen’s most abundant isotope



2a) 94 protons, 150 neutrons, 94 electrons;

2b) 8 protons, 8 neutrons, 10 electrons.

From the lecture series Chemistry and Our Universe
Taught by Professor Ron B. Davis, Jr., Ph.D