Transcription
When you think of the word “basketball”, what’s the first thing that comes to mind? An orange sphere, probably, or perhaps your favorite athlete driving home a dunk. Now, do you know what’s inside a basketball? Have you ever seen the inside of one? Even if you’ve never seen the inside of a basketball, you may have a general idea about what you would find.
The same thing is true for atoms; generations ago, people didn’t know exactly what was inside an atom, but they had some “ideas”. Around 400 BC, a Greek philosopher named Democritus came up with a theory that everything in the world was made of tiny indestructible particles called “atomos”, which means “uncuttable”. He believed that the properties of materials depended on the type of atomos they were composed of. For example, sour or sharp-tasting things were made of particles with pointy edges, sweet stuff was made of more rounded or smooth atoms, and metals were composed of hard atoms. Apart from this shape misinterpretation, he was basically on the right track about atomic composition dictating the properties of a substance. However, this theory was largely discredited by Aristotle—the original social influencer, who believed that everything on this planet was made of four elements: earth, fire, water, and air.
The next step in atomic theory development didn’t happen for nearly 2000 years, when British chemist and meteorology enthusiast John Dalton raised some interesting questions. He conducted experiments wherein he mixed two gases and observed their behavior. He noticed something rather curious when nitric oxide was allowed to interact with atmospheric oxygen. 36 measures of pure nitric acid reacted with 100 measures of air to create 80 measures of a new gas that was neither nitrous nor oxygenous. This intrigued him quite a bit, so he conducted the same experiment with different volumes of gas. He observed that the gases reacted with each other, but only in a fixed ratio. This gave rise to the law of multiple proportions and the theory of atomism. Following his breakthrough, Dalton proposed that everything in the world was made up of atoms—tiny indestructible solid spheres that were unique for every element. Atoms of different elements combine to form different compounds and are rearranged during chemical reactions. More than two centuries of scientific development have passed since he first proposed that idea, but some aspects of that model remain uncontested to this day.
Until the late 19th century, atoms were envisioned as indivisible particles, but that belief was shaken by an English physicist named J.J Thompson and his trusty cathode ray tube. Inside an almost vacuum glass tube, a visible beam of particles or cathode rays was generated by applying high voltage across metal electrodes. The stream of particles produced from the metal deflected away from the negative charge and directed more towards the positive charge. After repeating this experiment several times with other metals, he came up with the first atomic model—the famous “plum pudding” model. This model characterizes an atom as a particle that is composed of a positively charged mass (the pudding), as well as tiny negative charges embedded in it (like plums). After some initial resistance, this model became quite popular in the scientific world.
Even so, New Zealand-born Ernest Rutherford was not convinced. It was the early 1900s, radioactivity was all the rage, and during his work on radioactive decay, Rutherford discovered alpha, beta, and gamma rays. He wanted to develop a method to detect alpha particles and use it to probe into the structure of an atom. He did what every physicist at the time did… he came up with an experiment! The gold-foil experiment, also known as Geiger–Marsden experiments, consisted of a thin sheet of gold foil with a circular Zinc Sulphide-coated screen behind it. The screen would flash every time an alpha particle hit it. Rutherford expected the particles to bullet through the foil and hit the screen behind it. While most of the particles did behave as expected, some were deflected at an angle greater than 90 degrees. Backed by his observations, he came up with a new atomic model that disproved the previous one. He proposed an atomic structure where most of the atom’s mass was concentrated in a positively charged center (which he later named the nucleus) around which the electrons orbited like planets around the sun.
A year after the publication of Rutherford’s atomic theory, Niels Bohr found a discrepancy in the model. If electrons were orbiting around a positively charged center, at some point those electrons would lose their energy and collapse into the nucleus, thus making the atoms unstable. However, that wasn’t the case, as atoms were generally quite stable, aside from the radioactive ones. This is where quantum physics comes into the picture. He used the concept of quantized energy to propose that electrons moved around the nucleus in fixed orbits or shells. A shell closer to the nucleus has lower energy, while the one farthest away has the highest energy. If an electron jumped to a lower energy orbit, it would give out the extra energy in the form of radiation, thereby maintaining atomic stability. Even though Bohr’s model doesn’t hold true for complex multi-electron systems, this model is still the most popular representation of atomic structure in most textbooks.
No matter how much we try, there is no avoiding the complexities of quantum mechanics. With the establishment of the quantum behavior of entities like electrons, it became quite clear that Bohr’s atomic model didn’t satisfy the Heisenberg Uncertainty principle. According to this principle, it’s impossible to know the exact position and trajectory of electrons in an atom, which means they can’t exist in fixed orbits, as Bohr hypothesized. Combining the concept of wave-particle duality and the uncertainty principle, Erwin Schrödinger came up with the quantum mechanical model of an atom. In this model, the electrons did not revolve around the nucleus in circular orbits, but rather as electron clouds in an atomic orbital, a region inside the atom where the probability of finding an electron is the highest. He also formulated the Schrodinger-wave equations that would help us accurately calculate the energy levels of electrons in an atom. This new and improved atomic model does not tell us where an electron is, but where it could be.
Clearly, our understanding of “what’s inside an atom” has evolved remarkably over the last few centuries, but it was only possible because having a “General” idea about atomic structure simply wasn’t good enough for some people. They dug deeper, often dedicating their careers and lives to their pursuit, and now we know so much more about the stuff that makes up our planet. With that in mind, if you feel curious and adventurous someday and want to know what’s inside your basketball... you know what to do!