Evolution of Atomic Theory

Dalton's atomic theory laid important foundations in chemistry, but as time went on, parts of his theory came into question. Importantly, the assertion that atoms are indivisible seemed less likely. In 1896 French physicist Henri Becquerel discovered that some elements give off radiation similar to X-rays. Over the following two decades, Polish-born French physicist Marie Curie coined the term radioactive, which means spontaneously emitting energetic particles or radiation. She also discovered two new radioactive elements: radium and polonium.

The radioactive properties of elements led English physicist J.J. Thomson to use a cathode-ray tube to determine precisely what was happening. A cathode ray is a beam of electrons emitted from a negatively charged conducting plate in a vacuum chamber containing very little gas. Cathode rays can only be detected in a cathode-ray tube when they strike materials painted on the end of the tube called phosphors, materials that emit visible light when struck by electromagnetic radiation. A fluorescent substance is a type of phosphor that emits the visible light over a short time period.

Thomson wanted to know what exactly cathode rays were, if they were composed of anything at all. He knew that cathode rays traveled in straight lines within the tube. He conducted experiments in which he switched the gas in the tube, but he found that it had no effect on the cathode rays. Varying the metal that emitted the rays also seemed to have no effect. Thomson began to believe that the rays were composed of tiny particles that carried an electric charge. To demonstrate this idea, he placed positively and negatively charged plates alongside the beam in the tube. He found that the plates deflected the rays toward the positive side. This observation confirmed Thomson's suspicions and showed that cathode rays are negatively charged. Thomson called the negatively charged particles of cathode rays "corpuscles." We now know them as electrons. An electron is a negatively charged subatomic particle. Based on these experiments, Thomson proposed the plum pudding model of the atom in 1904, which shows negatively charged electrons floating in a positively charged material, similar to the way raisins are scattered throughout a British dish called plum pudding.

British physicist Ernest Rutherford was curious about Thomson's plum pudding model. Rutherford had isolated alpha particles. He had determined that an alpha particle ($\alpha$) is identical to a helium ion (He²⁺) that is emitted during the decay of radioactive elements. Rutherford also knew that α particles have significantly more mass than electrons. In 1909, using the plum pudding model, Rutherford hypothesized that if a beam of α particles is directed toward an atom, the particles would only be deflected if they hit an electron. This should happen rarely because electrons have very little mass and should be distributed randomly throughout the atom.

Rutherford's assistants Hans Geiger and Ernest Marsden conducted experiments to test the hypothesis. They set up a beam of alpha particles pointed directly toward a sheet of gold foil. They surrounded the gold foil with a detector that would emit light when struck by an alpha particle. Rutherford expected that most of the alpha particles would pass straight through the foil and strike the sheet behind it. Instead, Geiger and Marsden observed many alpha particles bouncing off the gold foil in multiple directions, including back toward the source emitting the beam. From these results, Rutherford proposed the nuclear model of the atom, in which an atom is composed of a positively charged center with electrons moving around it. Today, we know that Rutherford and his assistants had discovered the nucleus, which is the positively charged center of an atom containing protons and neutrons. At the time of Rutherford's experiments, neutrons had not yet been discovered. He proposed that the nucleus was composed of positively charged particles, which he named protons. A proton is a positively charged subatomic particle in the nucleus of an atom.

Niels Bohr, a Danish physicist, was also interested in discovering the nature of the atom. Rutherford's nuclear model had clarified the physical space of the atom, but classical physics predicts that electrons lose energy as they orbit the nucleus and should quickly fall into the nucleus. Yet, clearly, electrons remain in the area surrounding the nucleus. In 1913 Bohr suggested that electrons travel along specific paths around the nucleus. He called each path a shell and proposed that only shells with a specific, discrete radius from the nucleus are possible. Electrons cannot be found outside these shells. The radius r_n of each shell is represented as: The variable n is a positive integer representing the electron's state, its shell number, $n=1$, $n=2$, $n=3$, and so on, and r₁ is the Bohr radius, which is the distance between the nucleus and the electron in a hydrogen atom. Bohr calculated r₁ to be $0.529\times{10}^{-10}\;\rm{ m}$ for one-electron (hydrogen-like) systems. Bohr then turned his attention to electron energies. If electrons can only occupy shells at a particular distance from the nucleus, they should gravitate toward the lowest possible energy state. Bohr calculated this energy (E) for the nth level of hydrogen. This can be calculated by using the Rydberg formula, which is a mathematical formula of an electron moving between energy levels to predict the wavelength of light. The ground state energy level for a hydrogen atom is –13.6 eV. Thus, the lowest possible energy level, called the ground state, for a hydrogen electron is calculated by substituting 1 for n. This number is negative because it is relative to the energy of an electron that has been removed from its nucleus, which thus has an energy of 0 eV. An electron in orbit around a nucleus is more stable than an electron separated from its nucleus because it has less energy. Bohr's model only works with systems containing one electron. Bohr suggested that electrons in lower energy states (lower shells) can be excited to higher energy states (higher shells). However, electrons will always return to their lowest energy state. This return releases energy in the form of a photon, which is a packet of light. Bohr's model correctly explained and predicted absorption and emission spectra of elements. Thus, his model gained widespread acceptance and is commonly used today as an introduction to atomic theory.

Bohr's model gave the most accurate representation of the makeup of an atom at the time. Concurrently, beginning in 1909, American physicist Robert Millikan conducted a series of experiments to determine the charge on a single electron. He set up a device that would spray charged water droplets into an electric field. He measured their charges to be multiples of a single value, but imprecision in his experimental setup drew criticism of his results. He therefore replaced charged water droplets with charged oil droplets, which evaporated much more slowly, and obtained better results. He was thus able to calculate the charge of a single electron as approximately $1.6021766208\times{10}^{-19}\;\rm{C}$. In the years immediately after Bohr developed his atomic model, key discoveries were made that significantly changed scientists' understanding of the behavior of electrons in an atom. In 1926 Austrian physicist Erwin Schrödinger developed the quantum mechanical model of the atom in which he used mathematical equations to predict the positions of electrons. The main difference between the Bohr model and this model is the electron cloud, which is the arrangement of electrons moving around an atomic nucleus based on probabilities of their locations. Unlike the Bohr model, which describes the motion of each electron in a specific orbit, the quantum mechanical model describes an orbital, which is a region in which an electron has a high probability of being located. Orbitals are described by the quantum numbers s, p, d, and f, which differ from one another by their shapes.

Identifying the primary components of an atom was nearly complete. However, in 1932 English physicist James Chadwick began experiments to demonstrate the presence of another nuclear particle. Chadwick observed that, when struck by $\alpha$ particles, beryllium released radiation of an unknown kind. This radiation, in turn, caused other elements to release protons. Chadwick showed definitively that this radiation was not gamma rays because it was much more energetic than could be accounted for by gamma rays. By measuring the velocity of the protons released by the radiation, he determined the masses of the particles involved. In so doing, he discovered the neutron, a subatomic particle that has a neutral charge in the nucleus of an atom. The mass of a neutron is about the same as the mass of a proton.

With the discovery of the neutron, the identification of an atom's primary components was complete. All three subatomic particles, the proton, the neutron, and the electron, had been experimentally observed. The positions and charges of the three particles had been determined, along with their mass in the case of the more massive particles, the proton and the neutron.

The discovery of the neutron helped scientists understand why some elements had both radioactive and nonradioactive forms. Recall that an isotope is one of two or more forms of an element that have the same number of protons but different numbers of neutrons. The word isotope is often used to refer to the radioactive form, which tends to be the form with more neutrons than protons. For example, carbon has 6 protons, but carbon isotopes can have as few as 6 or as many as 17 neutrons. The isotope carbon-14 (¹⁴C), which has eight neutrons, is commonly used to determine the age of fossils through radiometric dating.

The discovery of subatomic particles and isotopes required modification of Dalton's original atomic theory. The first two parts of Dalton's theory have been partially disproven. Scientists now know that an atom is the smallest unit that has all the properties of an element, although it can be broken down further. Additionally, all atoms of the same element have the same number of protons, but their masses can vary slightly, depending on their number of neutrons. The atomic weight listed on the periodic table is the average mass of all isotopes of an element, based on the relative abundance of each isotope. The third and fourth parts of Dalton's theory hold true: two or more elements combine to form compounds, and chemical reactions are rearrangements of atoms without destruction or creation of new atoms.