The atom is defined as the smallest part of an element that retains the chemical properties of that element. The existence of atoms was first guessed as early as 400 BC, when Greek philosophers debated whether one could divide a substance into infinitely smaller pieces or if eventually a smallest, indivisible particle would be reached. Around 450 BC, the Greek philosopher Democritus proposed that all matter is made up of small, indivisible particles called atomos (meaning indivisible). His ideas were met with much criticism, and it was not until the early 1800s that the atomic theory took complete hold in scientific thought.
Atomic models have their beginnings in the early atomic theory. The law of conservation of mass, the law of definite proportions, and the law of multiple proportions provided models for the behavior of chemical reactions, but the laws could not be explained. In 1808, John Dalton (1766–1844) proposed his atomic theory, which served as an explanation for these phenomena. His theory consists of five postulates. The first postulate states that all matter is composed of atoms. Second, the atoms of a given element are all exactly the same in structure and properties. Third, atoms cannot be further divided, created, nor destroyed. Fourth, atoms of different elements combine to form compounds. Lastly, chemical reactions involve the combination, separation, and rearrangement of atoms. These five postulates not only explained the laws of conservation of mass, definite proportions, and multiple proportions, but also served as the basis for the study and development of various atomic models.
Not all of Dalton’s atomic theory has stood the test of time. The evolution of the atomic model has led to the discovery that atoms are divisible into smaller particles, and that not all atoms of a given element are exactly the same. In the late 1800s, scientists discovered that atoms are composed of three subatomic particles, called protons, neutrons, and electrons. The protons and neutrons are located in a central region of the atom called the nucleus, and the electrons occupy the space surrounding the nucleus. The number and arrangement of each of these particles within the atom determine the chemical properties of the element. Atoms of the same element can differ in the number of neutrons (called isotopes) and also in the number of electrons (called ions). The number of protons determines the chemical identity of the element. The discovery of these subatomic particles, along with Dalton’s atomic theory, set in motion the development of several atomic models. Advances in modern physics allowed scientists to find a large number of subatomic particles, including the quarks that are the fundamental subatomic particles that form protons and neutrons.
The discovery of the first subatomic particle, the electron, resulted from experiments involving the effects of electricity on matter. In the late 1800s, the cathode ray tube was developed and used in several investigations. A cathode ray tube is a partially evacuated glass tube containing a gas at low pressure. At one end of the tube is a cathode, at the other end, an anode. The cathode and anode are attached to a voltage source. The voltage source creates a current that can be passed through the gas trapped inside. Early experiments showed that the current caused the surface of the tube directly opposite the cathode to glow. It was hypothesized that a stream of particles originating at the cathode and moving toward the anode caused the glow. This stream of particles was called a cathode ray. When a paddle wheel was placed in the tube, it rolled from cathode to anode, which showed that the particles making up the cathode ray had mass. When exposed to a magnetic field, the cathode ray was deflected in the same manner as an electric current, which has a negative charge. Therefore, it was concluded that the particles that compose a cathode ray not only had mass but also a negative charge.
English physicist Joseph John Thomson (1856–1940) confirmed these findings in 1897. Thomson performed a series of experiments in which he was able to determine the ratio of the charge of the particles that make up the cathode ray to their mass by measuring the deflection of the rays with varying magnetic and electric fields. Thompson performed the same experiments using different metals for the cathode and anode as well as different gases inside the tube. His experiments demonstrated that the amount of deflection could be predicted mathematically. Thomson found that the charge-to-mass ratio was always the same, regardless of the materials used. He then concluded that all cathode rays are made up of the same particles, which were later named electrons by another English physicist, G. Johnstone Stoney (1826–1911).
The American physicist Robert A. Millikan (1868–1953) of the University of Chicago performed experiments that further confirmed Thomson’s results. Through his “oil droplet” experiments, in 1909 he discovered that the mass of one electron is approximately one two-thousandth that of a hydrogen atom. In the experiments, he used x rays to give oil droplets a negative charge. He then sprayed these droplets through an apparatus, allowing them to fall between two electrically charged plates. He varied the charge on the two plates and measured how this change affected the droplets’ rate of fall. Using these data, he calculated that the charge of every oil droplet was a multiple of the same number each time, and concluded that this must be the charge of a single electron. Using this number, and Thompson’s charge-to-mass ratio, he was able to calculate the mass of one electron. He reasoned that since cathode rays show the same deflection for any gas used, electrons must be present in the atoms of all elements. Since electrons have a negative charge, and all atoms are electrically neutral, there must also be a positive charge present in the atom. Further, because the mass of an electron is so much smaller than that of an entire atom, there must be other particles present in the atom to account for the majority of its mass. These results were the first to show that atoms are indeed divisible and were the basis for the first atomic model.
Thomson used these results to formulate his “plum pudding” model of the atom between the years 1903–1907. This model was an adaptation of a similar model first proposed by Lord Kelvin (1824–1907) in 1902. According to this model, the negatively charged electrons of an atom are found within a positively charged material, much like plums embedded in plum pudding. If one of the electrons were displaced, it would move back to its original position. This provided a stable model of a neutral atom. Around the same time period, the Japanese physicist Hantaro Nagoaka developed the “Saturnian” atomic model. In 1904, Nagaoka proposed that an atom resembled the planet Saturn. The planet itself was a region of positive charge around which rings of electrons circled. The atom, according to this model, was unstable because electrons moving in rings around a positive charge would gradually lose energy and eventually fall into the center region.
Thomson’s “plum pudding” model won favor over Nagaoka’s “Saturnian” model, but was accepted for only a few years. In 1911, New Zealand scientist Ernest Rutherford (1871–1937) proposed his own atomic model based on his famous gold foil experiments. Together with his colleagues Hans Geiger (1882–1945) and Ernest Mardsen, Rutherford aimed a stream of alpha particles at a thin sheet of gold foil. Alpha particles have a positive charge (+ 2) and are about four times as massive as a hydrogen atom. Their hypothesis was that the alpha particles would pass through the gold foil with minimal deflection, since mass and charge are distributed uniformly throughout an atom, as proposed by Thomson. The data did not agree with this assumption.
Some of the alpha particles were deflected at large angles as they passed through the gold foil. Even more surprising, about one in 8,000 particles was deflected straight back toward the source. As Rutherford described it, it was “as if you had fired a 15 in (38 cm) artillery shell at a piece of tissue paper and it came back and hit you.” He proposed that the deflected alpha particles must have come in contact with a densely packed positive charge. He called this region of positive charge the nucleus. The nucleus is surrounded by empty space, through which the electrons circled like the planets circle the sun. These experiments demonstrated that the atom has a tiny nucleus. Despite its minimal size, the nucleus contained most of the mass of the atom.
This model was not widely accepted by physicists because it was difficult to explain how such a small portion of the atom could carry most of the mass. It also suggested that the charge of the nucleus determined the properties of the atom, which disagreed with Dmitri Mendeleev’s periodic table of the elements. According to Mendeleev, the atomic mass of the element determined its properties, not the charge of the nucleus. Furthermore, it did not explain what kept the negatively charged electrons from falling into the positively charged nucleus.
The English scientist Henry Gwyn Jeffreys Moseley (1887–1915) soon solved the mystery of nuclear charge determining the properties of the atom. Moseley discovered in 1913 that each element contains a unique positive charge in its nucleus. The nucleus of the atom must contain positively charged particles, called protons. Therefore, the amount of charge, or the number of protons, in its nucleus determines an element’s identity. The number of protons in the nucleus is called the atomic number. Moseley asserted that the periodic table should be arranged in order of increasing atomic number instead of increasing atomic mass. This arrangement allowed for a complete periodic table, with correct predictions of where new elements were yet to be discovered.
The discovery of the proton resulted in another mystery. The mass of the hydrogen atom was known to be larger than the mass of a proton and an electron added together. Scientists searched for the source of the missing mass by assuming that another particle that also contributes to the mass of the atom must exist in the nucleus of the atom. This particle must be neutral in charge, since the positively charged protons cancel out the charge of the negatively charged electrons, and the atom as a whole is neutral. Because it is electrically neutral, the detection of this missing particle was problematic. Approximately 30 years after the electron was discovered, this third particle was found.
Irene Joliot-Curie (1897–1956) performed experiments in which she bombarded a beryllium sample with alpha particles. These experiments resulted in a new beam with a higher penetrating power than the beam of alpha particles. In 1932, the British scientist James Chadwick (1891–1974) discovered that this new beam was composed of particles of approximately the same mass as protons. In addition, magnetic or electric fields could not deflect this beam. Chadwick concluded that the beam must be made up of neutral particles of approximately the same size as protons, which he called neutrons. Neutrons, together with protons, make up the nucleus of the atom and contribute to the majority of its mass. Electrons are located in the empty space surrounding the atom, which makes up most of its volume. The mass of the electron is insignificant when compared to the mass of the protons and neutrons.
Even with the discovery of the proton, Rutherford’s atomic model still did not explain how electrons could have stable orbits around the nucleus. The development of a mathematical constant by the German physicist Max Planck (1858–1947) served as the basis for the next atomic model. Planck developed his constant in 1900 when explaining how light was emitted from hot objects. He hypothesized that electromagnetic radiation could only be associated with specific amounts of energy, which he called quanta. The energy lost or gained by an atom must occur in a quantum, which can be thought of as a “packet” containing a minimum amount of energy. He described the relationship between a quantum of energy and the frequency of the radiation emitted mathematically by the equation E=hλ (where E is the energy, in joules, of one quantum of radiation, and λ is the frequency of the radiation). The letter h symbolizes Planck’s constant.
In 1905 Albert Einstein (1879-1955) developed a theory stating that light has a dual nature. Light acts not only as a wave, but also as a particle. Each particle of light has a quantum of energy associated with it and is called a photon. The energy of a photon can be expressed using Planck’s equation. Einstein’s hypothesis helped explain the light emitted when current is passed through a gas in a cathode ray tube. An atom that has the lowest potential energy possible is said to be in the ground state. When a current passes through a gas at low pressure, the potential energy of the atoms increases. An atom having a higher potential energy than its ground state is said to be in an excited state. The excited state atom is unstable and will return to the ground state. When it does, it gives off the lost energy as electromagnetic radiation (a photon). When an electric current is passed through an elemental gas, a characteristic color of light is emitted. This light can be passed through a prism where it splits into various bands of light at specific wavelengths. These bands are known as the line-emission spectrum for that element. The line-emission spectrum for hydrogen was the first to be described mathematically. Scientists now faced the task of developing a model of the atom that could account for this mathematical relationship.
The Bohr model of the atom
In 1913, the Danish theorist Niels Bohr (1885–1962) developed his quantized shell model of the atom. Bohr modified Rutherford’s model by hypothesizing that the electrons orbit the nucleus in specific regions of fixed size and energy. The energy of the electron depends on the size of the orbit. Electrons in the smallest orbits have the least energy. An atom is stable when its electrons occupy orbits of the lowest possible energy. The energy of an electron increases as it occupies orbits farther and farther from the nucleus.
These orbits can be thought of as the rungs of a ladder. As a person climbs up a ladder, they step on one rung or another, but not in between rungs, because a person cannot stand on air. Likewise, the electrons of an atom can occupy one orbit or another, but cannot exist in between orbits. While in an orbit, the electron has a fixed amount of energy. The electron gains or loses energy by moving to a new orbit, either further from or closer to the nucleus.
When an electron falls from the excited state to the ground state, a photon is emitted with a specific energy. The energy of the photon is equal to the energy difference between the two orbits. The energy of each photon corresponds to a particular frequency of radiation given by Planck’s equation, E = hλ. Bohr was able to calculate the energy of the electron in a hydrogen atom by measuring the wavelengths of the light emitted in its line-emission spectrum. Bohr’s atomic model was very stable because the electron could not lose any more energy than it had in the smallest orbit. One major problem with Bohr’s model was that it could not explain the properties of atoms with more than one electron, and by the early 1920s, the search for a new atomic model had begun.
The development of quantum mechanics served as the foundation of the modern atomic theory. In 1922, the American physicist Arthur H. Compton (1892–1962) conduced x-ray scattering experiments that confirmed and advanced Einstein’s theory on the dual nature of light. In 1923, the French physicist Louis-Victor de Broglie (1892–1987) expanded on this theory by proposing that all matter, as well as radiation, behaves both as a particle and a wave. Until this time, scientists had viewed matter and energy as distinct phenomena that followed different laws.
Broglie’s proposal was not supported by experimental or mathematical evidence until 1926 when the Austrian physicist Erwin Schrödinger (1887–1961) developed his mathematical wave equation. Schrödinger proposed that electrons also behaved like waves. His wave equation could be used to find the frequency of the electrons and then Planck’s equation could be used to find the corresponding energy. Schrödinger’s equation gave a much more precise description of an electron’s location and energy than Bohr’s model could. It could also be used for atoms with more than one electron. Furthermore, only waves of specific frequencies could be solved using his equation. This demonstrated that only certain energies are possible for the electrons in an atom. Further experiments demonstrated that Broglie was correct in his assertion that matter could behave as waves, as electrons were diffracted and exhibited interference.
In 1927, German physicist Werner Heisenberg (1901–1976) developed what is now known as the Heisenberg uncertainty principle. This hypothesis states that the position and velocity of an electron, or any moving particle, cannot both be known at the same time. This meant that the solutions to the Schrödinger wave equation, known as wave functions, could describe only the probability of finding an electron in a given orbit. Therefore, the electrons are not located in discrete orbits, as hypothesized in the Bohr model, but instead occupy a hazier region, called an orbital. An orbital indicates a probable location of the electrons in an atom instead of a definite path that they follow. The probable location of the electrons in an orbital is described by a series of numbers called quantum numbers.
The quantum model of the atom uses four quantum numbers to describe the arrangement of electrons in an atom, much like an address describes the locations of houses on a street. This arrangement is known as the electron configuration. The atoms for each element have their own distinct electron configuration. The ground state electron configuration of an atom represents the lowest energy arrangement of the electrons in an atom. The placement of electrons in a particular configuration is based on three principles. The first, the Aufbau principle, states that an electron will occupy the lowest possible energy orbital available. The Pauli exclusion principle states that each electron in an atom has its own distinct set of four quantum numbers. No two electrons in an atom will have the same set.
Lastly, Hund’s rule states that even though each orbital can hold two electrons, the electrons will occupy the orbitals such that there are a maximum number of orbitals with only one electron. Developed by the German scientist Friedrich Hund (1896–1997), Hund’s rule allows scientists to predict the order in which electrons fill an atom’s suborbital shells.
Hund’s rule is based on the Aufbau principle, which states that electrons are added to the lowest available energy level of an atom. Around each atomic nucleus, electrons occupy energy levels termed shells. Each shell has a spherical s orbital, and, starting with the second shell, orbitals (p, d, f, etc.) and suborbitals (e.g., 2px, 2py, 2pz ) with differing size, shapes and orientation (i.e., direction in space).
Although each suborbital can hold two electrons, the electrons all carry negative charges, and, because like charges repel, electrons repel each other. In accord with Hund’s rule, electrons space themselves as far apart as possible by occupying all available vacant suborbitals before pairing up with another electron. The unpaired electrons all have the same spin quantum number (represented in electron configuration diagrams with arrows all pointing either upward or downward).
In accord with the Pauli exclusion principle, which states that each electron must have its own unique set of quantum numbers that specify its energy, and because all electrons have a spin of 1/2, each suborbital can hold up to two electrons only if their spins are paired +1/2 with–1/2. In electron configuration diagrams, paired electrons with opposite spins are represented by paired arrows pointing up and down.
Although Hund’s rule accurately predicts the electron configuration of most elements, exceptions exist, especially when atoms and ions have the opportunity to gain additional stability by having filled s shells or half-filled d or f orbitals.
In 1928, the English physicist P. A. M. Dirac (1902–1984) formulated a new equation to describe the electron. Schrödinger’s equation did not allow for the principles of relativity and could only be used to describe the movement of particles that are slower than the speed of light. Because electrons move at a much greater velocity, Dirac introduced four new wave functions to describe the behavior of electrons. These functions described electrons in various states. Two of the states corresponded to their spin orientations in the atom, but the other two could not be explained. In 1932, the American physicist Carl David Anderson (1905–1991) discovered the positron, which explained the two mystery states described by Dirac.
Alpha particle —Two protons and two neutrons bound together and emitted from the nucleus during some kinds of radioactive decay.
Anode —A positively charged electrode.
Cathode —A negatively charged electrode.
Electromagnetic radiation —The energy of photons, having properties of both particles and waves. The major wavelength bands are, from short to long: cosmic, ultraviolet, visible or “light,” infrared, and radio.
Law of conservation of mass —Mass is neither created nor destroyed during ordinary chemical or physical reactions.
Law of definite proportions —A chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or the source of the compound.
Law of multiple proportions —If two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers.
Positron —A positively charged particle having the same mass and magnitude of charge as the electron.
Potential energy —Stored energy.
Quarks —Believed to be the most fundamental units of protons and neutrons.
Subatomic particle —An elementary particle smaller than an atom. Protons, neutrons, and electrons are examples of subatomic particles.
Brown, Theodore E., et al. Chemistry: The Central Science. 10th ed. Upper Saddle River, NJ: Prentice Hall, 2005.
Cook, Norman D. Models of the Atomic Nucleus. New York: Springer, 2006.
Lodge, Oliver. Atoms and Rays: An Introduction to Modern Views on Atomic Structure and Radiation. Kila, MT: Kessinger Publishing, 2003.
Particle Data Group. Lawrence Berkeley National
Laboratory. “The Particle Adventure: The Fundamentals of Matter and Force” <http://particleadventure.org/particleadventure/> (accessed October 17, 2006).
K. Lee Lerner
"Atomic models." The Gale Encyclopedia of Science. . Encyclopedia.com. (August 8, 2019). https://www.encyclopedia.com/science/encyclopedias-almanacs-transcripts-and-maps/atomic-models
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