Atomic theory

"Atomic model" redirects here. For the unrelated term in mathematical logic, see atomic model (mathematical logic).

This article is about the historical development of understanding the existence and behavior of atoms. For a history of the study of how atoms combine to form molecules, see history of molecular theory. For the modern view of the atom which developed from atomic theory, see atomic physics.

The current theoretical model of the atom involves a dense nucleus surrounded by a probabilistic "cloud" of electrons

Atomic theory is the scientific theory that matter is composed of particles called atoms. The concept that matter is ultimately composed of discrete particles is an ancient idea, but gained scientific credence in the 19th century when scientists found it could explain the behavior of gases and how chemical elements reacted with each other. By the end of the 19th century, atomic theory had gained full acceptance in the scientific community. It was one of the most crucial scientific discoveries in history.

The term "atom" comes from the Greek word atomos, which means "uncuttable". John Dalton applied the term to the basic particles of the chemical elements under the mistaken belief that chemical atoms are the fundamental particles of matter. It was another century before scientists realized that Dalton's so-called atoms have an underlying structure of their own. Particles which are truly indivisible are now referred to as "elementary particles".

Philosophical atomism

Main article: Atomism

See also: Alchemy

The basic idea that matter is made up of tiny indivisible particles is an old idea that appeared in many ancient cultures. The word atom is derived from the ancient Greek word atomos, ^{[lower-alpha 1]} which means "uncuttable". This ancient idea was based in philosophical reasoning rather than scientific reasoning. Modern atomic theory is not based on these old concepts. ^{[1] [2]} In the early 19th century, the scientist John Dalton noticed that chemical substances seemed to combine with each other by discrete and consistent units of weight, and he decided to use the word atom to refer to these units. ^[3]

Groundwork

Near the end of the 18th century, a number of important developments in chemistry emerged without referring to the notion of an atomic theory. The first was Antoine Lavoisier redefining an element as a substance which scientists could not decompose into simpler substances by experimentation. This brought an end to the ancient idea of the elements of matter being fire, earth, air, and water, which had no experimental support. Lavoisier showed that water can be decomposed into hydrogen and oxygen, which in turn he could not decompose into anything simpler, thereby proving these are elements. ^[4] Lavoisier also defined the law of conservation of mass, which states that in a chemical reaction, matter does not appear nor disappear into thin air; the total mass remains the same even if the substances involved were transformed. Finally, there was the law of definite proportions, established by the French chemist Joseph Proust in 1797, which states that if a compound is broken down into its constituent chemical elements, then the masses of those constituents will always have the same proportions by weight, regardless of the quantity or source of the original compound. ^[5]

Dalton's law of multiple proportions

From A New System of Chemical Philosophy (John Dalton 1808).

John Dalton studied data gathered by himself and by other scientists. He noticed a pattern that later came to be known as the law of multiple proportions: in compounds which contain two particular elements, the amount of Element A per measure of Element B will differ across these compounds by ratios of small whole numbers. This suggested that each element combines with other elements in multiples of a basic quantity.

In 1804, Dalton explained his atomic theory to his friend and fellow chemist Thomas Thomson, who published an explanation of Dalton's theory in his book A System of Chemistry in 1807. According to Thomson, Dalton's idea first occurred to him when experimenting with "olefiant gas" (ethylene) and "carburetted hydrogen gas" (methane). Dalton found that "carburetted hydrogen gas" contains twice as much hydrogen per measure of carbon as "olefiant gas", and concluded that a molecule of "olefiant gas" is one carbon atom and one hydrogen atom, and a molecule of "carburetted hydrogen gas" is one carbon atom and two hydrogen atoms. In reality, an ethylene molecule has two carbon atoms and four hydrogen atoms (C₂H₄), and a methane molecule has one carbon atom and four hydrogen atoms (CH₄). ^[6] In this particular case, Dalton was mistaken about the formulas of these compounds, and it wasn't his only mistake. But in other cases, he got their formulas right. The following examples come from Dalton's own books A New System of Chemical Philosophy (in two volumes, 1808 and 1817):

Example 1 — tin oxides: Dalton identified two types of tin oxide. One is a grey powder that Dalton referred to as "the protoxide of tin", which is 88.1% tin and 11.9% oxygen. The other is a white powder which Dalton referred to as "the deutoxide of tin", which is 78.7% tin and 21.3% oxygen. Adjusting these figures, in the grey powder there is about 13.5 g of oxygen for every 100 g of tin, and in the white powder there is about 27 g of oxygen for every 100 g of tin. 13.5 and 27 form a ratio of 1:2. These compounds are known today tin(II) oxide (SnO) and tin(IV) oxide (SnO₂). In Dalton's terminology, a "protoxide" is a molecule containing a single oxygen atom, and a "deutoxide" molecule has two. ^{[7] [8]} (Tin oxides are actually crystals, they don't exist in molecular form.)

Example 2 — iron oxides: Dalton identified two oxides of iron. There is one type of iron oxide that is a black powder which Dalton referred to as "the protoxide of iron", which is 78.1% iron and 21.9% oxygen. The other iron oxide is a red powder, which Dalton referred to as "the intermediate or red oxide of iron" which is 70.4% iron and 29.6% oxygen. Adjusting these figures, in the black powder there is about 28 g of oxygen for every 100 g of iron, and in the red powder there is about 42 g of oxygen for every 100 g of iron. 28 and 42 form a ratio of 2:3. These compounds are iron(II) oxide and iron(III) oxide and their formulas are Fe₂O₂ and Fe₂O₃ respectively (iron(II) oxide's formula is normally written as FeO, but here it is written as Fe₂O₂ to contrast it with the other oxide). Dalton described the "intermediate oxide" as being "2 atoms protoxide and 1 of oxygen", which adds up to two atoms of iron and three of oxygen. That averages to one and a half atoms of oxygen for every iron atom, putting it midway between a "protoxide" and a "deutoxide". ^{[9] [10]}

Example 3 — nitrogen oxides: Dalton was aware of three oxides of nitrogen: "nitrous oxide", "nitrous gas", and "nitric acid". ^[11] These compounds are known today as nitrous oxide, nitric oxide, and nitrogen dioxide respectively. "Nitrous oxide" is 63.3% nitrogen and 36.7% oxygen, which means it has 80 g of oxygen for every 140 g of nitrogen. "Nitrous gas" is 44.05% nitrogen and 55.95% oxygen, which means there are 160 g of oxygen for every 140 g of nitrogen. "Nitric acid" is 29.5% nitrogen and 70.5% oxygen, which means it has 320 g of oxygen for every 140 g of nitrogen. 80 g, 160 g, and 320 g form a ratio of 1:2:4. The formulas for these compounds are N₂O, NO, and NO₂. ^{[12] [13]}

Dalton defined an atom as being the "ultimate particle" of a chemical substance, and he used the term "compound atom" to refer to "ultimate particles" which contain two or more elements. This is inconsistent with the modern definition, wherein an atom is the basic particle of a chemical element and a molecule is an agglomeration of atoms. The term "compound atom" was confusing to some of Dalton's contemporaries as the word "atom" implies indivisibility, but he responded that if a carbon dioxide "atom" is divided, it ceases to be carbon dioxide. The carbon dioxide "atom" is indivisible in the sense that it cannot be divided into smaller carbon dioxide particles. ^{[3] [14]}

Dalton made the following assumptions on how "elementary atoms" combined to form "compound atoms" (what we today refer to as molecules). When two elements can only form one compound, he assumed it was one atom of each, which he called a "binary compound". If two elements can form two compounds, the first compound is a binary compound and the second is a "ternary compound" consisting of one atom of the first element and two of the second. If two elements can form three compounds between them, then the third compound is a "quaternary" compound containing one atom of the first element and three of the second. ^[15] Dalton thought that water was a "binary compound", i.e. one hydrogen atom and one oxygen atom. Dalton did not know that in their natural gaseous state, the ultimate particles of oxygen, nitrogen, and hydrogen exist in pairs (O₂, N₂, and H₂). Nor was he aware of valencies. These properties of atoms were discovered later in the 19th century.

Because atoms are too small to be directly weighed using the methods of the 19th century, Dalton instead expressed the weights of the myriad atoms as multiples of the hydrogen atom's weight, which Dalton knew was the lightest element. By his measurements, 7 grams of oxygen will combine with 1 gram of hydrogen to make 8 grams of water with nothing left over, and assuming a water molecule to be one oxygen atom and one hydrogen atom, he concluded that oxygen's atomic weight is 7. In reality it is 16. Aside from the crudity of early 19th century measurement tools, the main reason for this error was that Dalton didn't know that the water molecule in fact has two hydrogen atoms, not one. Had he known, he would have doubled his estimate to a more accurate 14. This error was corrected in 1811 by Amedeo Avogadro. Avogadro proposed that equal volumes of any two gases, at equal temperature and pressure, contain equal numbers of molecules (in other words, the mass of a gas's particles does not affect the volume that it occupies). ^[16] Avogadro's hypothesis, now usually called Avogadro's law, provided a method for deducing the relative weights of the molecules of gaseous elements, for if the hypothesis is correct relative gas densities directly indicate the relative weights of the particles that compose the gases. This way of thinking led directly to a second hypothesis: the particles of certain elemental gases were pairs of atoms, and when reacting chemically these molecules often split in two. For instance, the fact that two liters of hydrogen will react with just one liter of oxygen to produce two liters of water vapor (at constant pressure and temperature) suggested that a single oxygen molecule splits in two in order to form two molecules of water. The formula of water is H₂O, not HO. Avogadro measured oxygen's atomic weight to be 15.074. ^[17]

Opposition to atomic theory

Dalton's atomic theory was not immediately accepted by all scientists. The law of multiple proportions was shown to not be a universal law when it came to organic substances. For instance, in oleic acid there is 34 g of hydrogen for every 216 g of carbon, and in methane there is 72 g of hydrogen for every 216 g of carbon. 34 and 72 form a ratio of 17:36, which is not a ratio of small whole numbers. We know now that carbon-based substances can have very large molecules, larger than any the other elements can form. Oleic acid's formula is C₁₈H₃₄O₂ and methane's is CH₄. ^[18] The law of multiple proportions by itself was not complete proof, and it wouldn't be until the end of the 19th century that atomic theory was universally accepted.

One problem was the lack of uniform nomenclature. The word "atom" implied indivisibility, but Dalton defined an atom as being the ultimate particle of any chemical substance, not just the elements or even matter per se. This meant that "compound atoms" such as carbon dioxide could be divided, as opposed to "elementary atoms". Dalton disliked the word "molecule", regarding it as "diminutive". ^{[3] [19]} Amadeo Avogadro did the opposite, he exclusively used the word "molecule" in his writings, eschewing the word "atom", instead using the term "elementary molecule". ^[20] Jöns Jacob Berzelius used the term "organic atoms" to refer to particles containing three or more elements, because he thought this only existed in organic compounds. Jean-Baptiste Dumas used the terms "physical atoms" and "chemical atoms"; a "physical atom" is a particle that cannot be divided by physical forces such as temperature and pressure, and a "chemical atom" is a particle that cannot be divided by chemical reactions. ^[21]

The modern definitions of atom and molecule—an atom being the basic particle of an element, and a molecule being an agglomeration of atoms—were established in the late half of the 19th century. A key event was the Karlsruhe Congress in Germany in 1860. As the first international congress of chemists, its goal was to establish some standards in the community. A major proponent of the modern distinction between atoms and molecules was Stanislao Cannizzaro.

The various quantities of a particular element involved in the constitution of different molecules are integral multiples of a fundamental quantity that always manifests itself as an indivisible entity and which must properly be named atom.

—  Stanislao Cannizzaro, 1860 ^[22]

Cannizzaro criticized past chemists such as Berzelius for not accepting that the particles of certain gaseous elements are actually pairs of atoms, which led to mistakes in their formulation of certain compounds. Berzelius believed that hydrogen gas and chlorine gas particles are solitary atoms. But he observed that when one liter of hydrogen reacts with one liter of chlorine, they form two liters of hydrogen chloride instead of one. Berzelius decided that Avogadro's law does not apply to compounds. Cannizzaro preached that if scientists just accepted the existence of single-element molecules, such discrepancies in their findings would be easily resolved. But Berzelius did not even have a word for that. Berzelius used the term "elementary atom" for a gas particle which contained just one element and "compound atom" for particles which contained two or more elements, but there was nothing to distinguish H₂ from H since Berzelius didn't believe in H₂. So Cannizzaro called for a redefinition so that scientists could understand that a hydrogen molecule can split into two atoms in the course of a chemical reaction. ^[23]

A second objection to atomic theory was philosophical. Scientists in the 19th century had no way of directly observing atoms. They inferred the existence of atoms through indirect observations, such as Dalton's law of multiple proportions. Some scientists, notably those who ascribed to the school of positivism, argued that scientists should not attempt to deduce the deeper reality of the universe but only systemize what patterns they can directly observe. The anti-atomists argued that while atoms might be a useful abstraction for predicting how elements react, they do not reflect concrete reality.

Such scientists were sometimes known as "equivalentists", because they preferred the theory of equivalent weights, which is a generalization of Proust's law of definite proportions. For example, 1 gram of hydrogen will combine with 8 grams of oxygen to form 9 grams of water, therefore the "equivalent weight" of oxygen is 8 grams. This position was eventually quashed by two important advancements that happened later in the 19th century: the development of the periodic table and the discovery that molecules have an internal architecture that determines their properties. ^[24]

Isomerism

Scientists discovered some substances have the exact same chemical content but different properties. For instance, in 1827, Friedrich Wöhler discovered that silver fulminate and silver cyanate are both 107 parts silver, 12 parts carbon, 14 parts nitrogen, and 16 parts oxygen (we now know their formulas as both AgCNO). In 1830 Jöns Jacob Berzelius introduced the term isomerism to describe the phenomenon. In 1860, Louis Pasteur hypothesized that the molecules of isomers might have the same set of atoms but in different arrangements. ^[25]

In 1874, Jacobus Henricus van 't Hoff proposed that the carbon atom bonds to other atoms in a tetrahedral arrangement. Working from this, he explained the structures of organic molecules in such a way that he could predict how many isomers a compound could have. Consider, for example, pentane (C₅H₁₂). In van 't Hoff's way of modelling molecules, there are three possible configurations for pentane, and scientists did go on to discover three and only three isomers of pentane. ^{[26] [27]}

n-pentane

isopentane

neopentane

Jacobus Henricus van 't Hoff's way of modelling molecular structures correctly predicted the three isomers of pentane (C₅H₁₂).

Isomerism was not something that could be fully explained by alternative theories to atomic theory, such as radical theory and the theory of types. ^{[28] [29]}

Mendeleev's periodic table

Main article: Periodic table

Dmitrii Mendeleev noticed that when he arranged the elements in a row according to their atomic weights, there was a certain periodicity to them. ^{[30] : 117} For instance, the second element, lithium, had similar properties to the ninth element, sodium, and the sixteenth element, potassium — a period of seven. Likewise, beryllium, magnesium, and calcium were similar and all were seven places apart from each other on Mendeleev's table. Using these patterns, Mendeleev predicted the existence and properties of new elements, which were later discovered in nature: scandium, gallium, and germanium. ^{[30] : 118} Moreover, the periodic table could predict how many atoms of other elements that an atom could bond with — e.g., germanium and carbon are in the same group on the table and their atoms both combine with two oxygen atoms each (GeO₂ and CO₂). Mendeleev found these patterns validated atomic theory because it showed that the elements could be categorized by their atomic weight. Inserting a new element into the middle of a period would break the parallel between that period and the next, and would also violate Dalton's law of multiple proportions. ^[31]

Mendeleev's periodic table from 1871.

In the modern periodic table, the periodicity of the elements mentioned above is eight rather than seven because the noble gases were not known back when Mendeleev devised his table. The rows also now have different lengths (2, 8, 18, and 32) to fit with quantum theory.

Brownian motion

In 1827, the British botanist Robert Brown observed that dust particles inside pollen grains floating in water constantly jiggled about for no apparent reason. In 1905, Albert Einstein theorized that this Brownian motion was caused by the water molecules continuously knocking the grains about, and developed a mathematical model to describe it. This model was validated experimentally in 1908 by French physicist Jean Perrin, who used Einstein's equations to measure the size of atoms. ^[32]

Kinetic diameters of various simple molecules

Molecule	Perrin's measurements [33]	Modern measurements
Helium	1.7 × 10−10 m	2.6 × 10−10 m
Argon	2.7 × 10−10 m	3.4 × 10−10 m
Mercury	2.8 × 10−10 m	3 × 10−10 m
Hydrogen	2 × 10−10 m	2.89 × 10−10 m
Oxygen	2.6 × 10−10 m	3.46 × 10−10 m
Nitrogen	2.7 × 10−10 m	3.64 × 10−10 m
Chlorine	4 × 10−10 m	3.20 × 10−10 m

Statistical mechanics

In order to introduce the Ideal gas law and statistical forms of physics, it was necessary to postulate the existence of atoms. In 1738, Swiss physicist and mathematician Daniel Bernoulli postulated that the pressure of gases and heat were both caused by the underlying motion of molecules.

In 1860, James Clerk Maxwell, who was a vocal proponent of atomism, was the first to use statistical mechanics in physics. ^[34] Ludwig Boltzmann and Rudolf Clausius expanded his work on gases and the laws of thermodynamics especially the second law relating to entropy. In the 1870s, Josiah Willard Gibbs extended the laws of entropy and thermodynamics and coined the term "statistical mechanics." Einstein later independently reinvented Gibbs' laws, because they had only been printed in an obscure American journal. ^[35] Einstein later commented that had he known of Gibbs' work, he would "not have published those papers at all, but confined myself to the treatment of some few points [that were distinct]." ^[36] All of statistical mechanics and the laws of heat, gas, and entropy took the existence of atoms as a necessary postulate.

Discovery of the electron

Main articles: Electron and Plum pudding model

The cathode rays (blue) were emitted from the cathode, sharpened to a beam by the slits, then deflected as they passed between the two electrified plates.

Atoms were thought to be the smallest possible division of matter until 1897 when J. J. Thomson discovered the electron through his work on cathode rays. ^[37]

A Crookes tube is a sealed glass container in which two electrodes are separated by a vacuum. When a voltage is applied across the electrodes, cathode rays are generated, creating a glowing patch where they strike the glass at the opposite end of the tube. Through experimentation, Thomson discovered that the rays could be deflected by electric fields and magnetic fields. He concluded that these rays, rather than being a form of light, were composed of very light negatively charged particles. Thomson called these "corpuscles", but other scientists called them electrons, following an 1894 suggestion by George Johnstone Stoney for naming the basic unit of electrical charge. ^[38] He measured the mass of the electron to be 1,800 times smaller than that of hydrogen, the smallest atom. Electrons are so light yet carry so much charge that Thomson concluded they must be the basic particles of electrical effects. ^[39]

Thomson was trying to explain electricity but his theory entailed that the atom is not the indivisible entity scientists thought it was. A negatively charged part of the atom splits off to join with an adjacent atom in the action of an electrical current. When electrons do not flow their negative charge must balanced out by some source of positive charge within the atom so as to render the atom electrically neutral. This led to what became known as the plum pudding model, which imagined the electrons embedded in a sphere of positive charge like bits of fruit in a Christmas pudding. ^[40]

Discovery of the nucleus

Main article: Geiger–Marsden experiments

The Geiger–Marsden experiment
Left: Expected results: alpha particles passing through the plum pudding model of the atom with negligible deflection.
Right: Observed results: a small portion of the particles were deflected by the concentrated positive charge of the nucleus.

Thomson's plum pudding model was supplanted in 1909 by one of his former students, Ernest Rutherford, who discovered that most of the mass and positive charge of an atom is concentrated in a very small fraction of its volume, which he assumed to be at the very center.

Ernest Rutherford and his colleagues Hans Geiger and Ernest Marsden came to have doubts about the Thomson model after they encountered difficulties when they tried to build an instrument to measure the charge-to-mass ratio of alpha particles (these are positively-charged particles emitted by certain radioactive substances such as radium). The alpha particles were being scattered by the air in the detection chamber, which made the measurements unreliable. Thomson had encountered a similar problem in his work on cathode rays, which he solved by creating a near-perfect vacuum in his instruments. Rutherford didn't think he'd run into this same problem because alpha particles are much heavier than electrons. According to Thomson's model of the atom, the positive charge in the atom is not concentrated enough to produce an electric field strong enough to deflect an alpha particle, and the electrons are so lightweight they should be pushed aside effortlessly by the much heavier alpha particles. Yet there was scattering, so Rutherford and his colleagues decided to investigate this scattering carefully. ^[41]

Between 1908 and 1913, Rutherford and his colleagues performed a series of experiments in which they bombarded thin foils of metal with alpha particles. They spotted alpha particles being deflected by angles greater than 90°. To explain this, Rutherford proposed that the positive charge of the atom is not distributed throughout the atom's volume as Thomson believed, but is concentrated in a tiny nucleus at the center. Only such an intense concentration of charge could produce an electric field strong enough to deflect the alpha particles as observed. ^[41] Rutherford's model is sometimes called the "planetary model". ^[42]

The Bohr model

Main article: Bohr model

The planetary model of the atom had two significant shortcomings. The first is that, unlike planets orbiting a sun, electrons are charged particles. An accelerating electric charge is known to emit electromagnetic waves according to the Larmor formula in classical electromagnetism. An orbiting charge should steadily lose energy and spiral toward the nucleus, colliding with it in a small fraction of a second. The second problem was that the planetary model could not explain the highly peaked emission and absorption spectra of atoms that were observed.

The Bohr model of the atom

Quantum theory revolutionized physics at the beginning of the 20th century, when Max Planck and Albert Einstein postulated that light energy is emitted or absorbed in discrete amounts known as quanta (singular, quantum). This led to a series of quantum atomic models such as the quantum model of Arthur Erich Haas in 1910 and the 1912 John William Nicholson quantum atomic model that quantized angular momentum as h/2π. ^{[43] [44]} In 1913, Niels Bohr incorporated this idea into his Bohr model of the atom, in which an electron could only orbit the nucleus in particular circular orbits with fixed angular momentum and energy, its distance from the nucleus (i.e., their radii) being proportional to its energy. ^[45] Under this model an electron could not spiral into the nucleus because it could not lose energy in a continuous manner; instead, it could only make instantaneous "quantum leaps" between the fixed energy levels. ^[45] When this occurred, light was emitted or absorbed at a frequency proportional to the change in energy (hence the absorption and emission of light in discrete spectra). ^[45]

Bohr's model was not perfect. It could only predict the spectral lines of hydrogen, not those of multielectron atoms. ^[46] Worse still, it could not even account for all features of the hydrogen spectrum: as spectrographic technology improved, it was discovered that applying a magnetic field caused spectral lines to multiply in a way that Bohr's model couldn't explain. In 1916, Arnold Sommerfeld added elliptical orbits to the Bohr model to explain the extra emission lines, but this made the model very difficult to use, and it still couldn't explain more complex atoms. ^{[47] [48]}

Discovery of isotopes

Main article: Isotope

While experimenting with the products of radioactive decay, in 1913 radiochemist Frederick Soddy discovered that there appeared to be more than one variety of some elements. ^[49] The term isotope was coined by Margaret Todd as a suitable name for these varieties. ^[50]

That same year, J. J. Thomson conducted an experiment in which he channeled a stream of neon ions through magnetic and electric fields, striking a photographic plate at the other end. He observed two glowing patches on the plate, which suggested two different deflection trajectories. Thomson concluded this was because some of the neon ions had a different mass. ^[51] The nature of this differing mass would later be explained by the discovery of neutrons in 1932: all atoms of the same element contain the same number of protons, while different isotopes have different numbers of neutrons. ^[52]

Discovery of nuclear particles

Main articles: Atomic nucleus and Discovery of the neutron

In 1917 Rutherford bombarded nitrogen gas with alpha particles and observed hydrogen nuclei being emitted from the gas (Rutherford recognized these, because he had previously obtained them bombarding hydrogen with alpha particles, and observing hydrogen nuclei in the products). Rutherford concluded that the hydrogen nuclei emerged from the nuclei of the nitrogen atoms themselves (in effect, he had split a nitrogen). ^[53]

From his own work and the work of his students Bohr and Henry Moseley, Rutherford knew that the positive charge of any atom could always be equated to that of an integer number of hydrogen nuclei. This, coupled with the atomic mass of many elements being roughly equivalent to an integer number of hydrogen atoms – then assumed to be the lightest particles – led him to conclude that hydrogen nuclei were singular particles and a basic constituent of all atomic nuclei. He named such particles protons. Further experimentation by Rutherford found that the nuclear mass of most atoms exceeded that of the protons it possessed; he speculated that this surplus mass was composed of previously unknown neutrally charged particles, which were tentatively dubbed "neutrons".

In 1928, Walter Bothe observed that beryllium emitted a highly penetrating, electrically neutral radiation when bombarded with alpha particles. It was later discovered that this radiation could knock hydrogen atoms out of paraffin wax. Initially it was thought to be high-energy gamma radiation, since gamma radiation had a similar effect on electrons in metals, but James Chadwick found that the ionization effect was too strong for it to be due to electromagnetic radiation, so long as energy and momentum were conserved in the interaction. In 1932, Chadwick exposed various elements, such as hydrogen and nitrogen, to the mysterious "beryllium radiation", and by measuring the energies of the recoiling charged particles, he deduced that the radiation was actually composed of electrically neutral particles which could not be massless like the gamma ray, but instead were required to have a mass similar to that of a proton. Chadwick now claimed these particles as Rutherford's neutrons. ^[54] For his discovery of the neutron, Chadwick received the Nobel Prize in 1935. ^[55]

Quantum mechanical models

Main article: History of quantum mechanics

The five filled atomic orbitals of a neon atom separated and arranged in order of increasing energy from left to right, with the last three orbitals being equal in energy. Each orbital holds up to two electrons, which most probably exist in the zones represented by the colored bubbles. Each electron is equally present in both orbital zones, shown here by color only to highlight the different wave phase.

In 1924, Louis de Broglie proposed that all particles—particularly subatomic particles such as electrons—have an associated wave. Erwin Schrödinger, fascinated by this idea, developed an equation ^[56] that describes an electron as a wave function instead of a point. This approach predicted many of the spectral phenomena that Bohr's model failed to explain, but it was difficult to visualize, and faced opposition. ^[57] One of its critics, Max Born, proposed instead that Schrödinger's wave function did not describe the physical extent of an electron (like a charge distribution in classical electromagnetism), but rather gave the probability that an electron would, when measured, be found at a particular point. ^[58] This reconciled the ideas of wave-like and particle-like electrons: the behavior of an electron, or of any other subatomic entity, has both wave-like and particle-like aspects, and whether one aspect or the other is observed depend upon the experiment. ^[59]

Schrödinger's wave model for hydrogen replaced Bohr's model, with its neat, clearly defined circular orbits. The modern model of the atom describes the positions of electrons in an atom in terms of probabilities. An electron can potentially be found at any distance from the nucleus, but, depending on its energy level and angular momentum, exists more frequently in certain regions around the nucleus than others; this pattern is referred to as its atomic orbital. The orbitals come in a variety of shapes—sphere, dumbbell, torus, etc.—with the nucleus in the middle. ^[60] The shapes of atomic orbitals are found by solving the Schrödinger equation. ^[61] Analytic solutions of the Schrödinger equation are known for very few relatively simple model Hamiltonians including the hydrogen atom and the hydrogen molecular ion. ^[62] Beginning with helium atom—which contains just two electrons—numerical methods are used to solve the Schrödinger equation. ^[63]

Qualitatively the shape of the atomic orbitals of multi-electron atoms resemble the states of the hydrogen atom. The Pauli principle requires the distribution of these electrons within the atomic orbitals such that no more than two electrons are assigned to any one orbital; this requirement profoundly affects the atomic properties and ultimately the bonding of atoms in to molecules. ^{[64] : 182}

See also

• Physics portal

• Spectroscopy
• History of molecular theory
• Timeline of chemical element discoveries
• Introduction to quantum mechanics
• Kinetic theory of gases
• Atomism
• The Physical Principles of the Quantum Theory

Footnotes

1. a combination of the negative term "a-" and "τομή," the term for "cut"

1. Pullman, Bernard (1998). The Atom in the History of Human Thought. Oxford, England: Oxford University Press. pp. 31–33. ISBN   978-0-19-515040-7. Archived from the original on 5 February 2021. Retrieved 25 October 2020.
2. Melsen (1952). From Atomos to Atom, pp. 18–19
3. Pullman (1998). The Atom in the History of Human Thought, p. 201
4. Pullman (1998). The Atom in the History of Human Thought. p. 197
5. "Law of definite proportions | chemistry". Encyclopedia Britannica. Retrieved 2020-09-03.
6. Thomas Thomson (1831). A History of Chemistry, Volume 2. p. 291
7. Dalton (1817). A New System of Chemical Philosophy vol. 2, p. 36
8. Melsen (1952). From Atomos to Atom. p. 137
9. Dalton (1817). A New System of Chemical Philosophy vol. 2. pp. 28-34: "the intermediate or red oxide is 2 atoms protoxide and 1 of oxygen"
10. Millington (1906). John Dalton, p. 113
11. Dalton (1808). A New System of Chemical Philosophy vol. 1, pp. 316–319
12. Dalton (1808). A New System of Chemical Philosophy vol. 1. pp. 316–319
13. Holbrow et al. (2010). Modern Introductory Physics, pp. 65–66
14. Dalton, quoted in Freund (1904). The Study of Chemical Composition. p. 288: "I have chosen the word atom to signify these ultimate particles in preference to particle, molecule, or any other diminiutive term, because I conceive it is much more expressive; it includes in itself the notion of indivisible, which the other terms do not. It may, perhaps, be said that I extend the application of it too far when I speak of compound atoms; for instance, I call an ultimate particle of carbonic acid a compound atom. Now, though this atom may be divided, yet it ceases to become carbonic acid, being resolved by such division into charcoal and oxygen. Hence I conceive there is no inconsistency in speaking of compound atoms and that my meaning cannot be misunderstood."
15. Dalton (1817). A New System of Chemical Philosophy vol. 1, pp. 213–214
16. Avogadro, Amedeo (1811). "Essay on a Manner of Determining the Relative Masses of the Elementary Molecules of Bodies, and the Proportions in Which They Enter into These Compounds". Journal de Physique. 73: 58–76.
17. Avogadro, Amedeo (1811). "Essai d'une manière de déterminer les masses relatives des molécules élémentaires des corps, et les proportions selon lesquelles elles entrent dans ces combinaisons". Journal de Physique. 73: 58–76. English translation
18. Trusted (1999). The Mystery of Matter, p. 73
19. Freund (1904). The Study of Chemical Composition. p. 288
20. Pullman (1998). The Atom in the History of Human Thought, p. 202
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Bibliography

• Andrew G. van Melsen (1960) [First published 1952]. From Atomos to Atom: The History of the Concept Atom. Translated by Henry J. Koren. Dover Publications. ISBN   0-486-49584-1.
• J. P. Millington (1906). John Dalton. J. M. Dent & Co. (London); E. P. Dutton & Co. (New York).
• Jaume Navarro (2012). A History of the Electron: J. J. and G. P. Thomson. Cambridge University Press. ISBN   978-1-107-00522-8.
• Jennifer Trusted (1999). The Mystery of Matter. MacMillan. ISBN   0-333-76002-6.
• Bernard Pullman (1998). The Atom in the History of Human Thought. Translated by Axel Reisinger. Oxford University Press. ISBN   0-19-511447-7.
• Jean Perrin (1910) [1909]. Brownian Movement and Molecular Reality. Translated by F. Soddy. Taylor and Francis.
• Ida Freund (1904). The Study of Chemical Composition. Cambridge University Press.
• Thomas Thomson (1807). A System of Chemistry: In Five Volumes, Volume 3. John Brown.
• Thomas Thomson (1831). The History of Chemistry, Volume 2. H. Colburn, and R. Bentley.
• John Dalton (1808). A New System of Chemical Philosophy vol. 1.
• John Dalton (1817). A New System of Chemical Philosophy vol. 2.
• Stanislao Cannizzaro (1858). Sketch of a Course of Chemical Philosophy. The Alembic Club.

Further reading

• Charles Adolphe Wurtz (1881) The Atomic Theory, D. Appleton and Company, New York.
• Alan J. Rocke (1984) Chemical Atomism in the Nineteenth Century: From Dalton to Cannizzaro, Ohio State University Press, Columbus (open access full text at http://digital.case.edu/islandora/object/ksl%3Ax633gj985).

External links

• Atomism by S. Mark Cohen.
• Atomic Theory – detailed information on atomic theory with respect to electrons and electricity.
• The Feynman Lectures on Physics Vol. I Ch. 1: Atoms in Motion