Atomic theory

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Various atoms and molecules as depicted in John Dalton's A New System of Chemical Philosophy (1808).

In chemistry and physics, atomic theory is a theory of the nature of matter, which states that matter is composed of discrete units called atoms, as opposed to antiquated beliefs that matter could be divided into any arbitrarily small quantity.

Atomic theory began thousands of years ago as a philosophical concept, and in the 19th century achieved widespread scientific acceptance thanks to discoveries in the field of stoichiometry. The chemists of the era believed the basic units of the elements were also the fundamental particles of nature and named them atoms (derived from the Greek word atomos, meaning "indivisible"). However, around the turn of the century, through various experiments with electromagnetism and radioactivity, physicists discovered that the so-called "indivisible atom" was actually a conglomerate of various subatomic particles (chiefly, electrons, protons and neutrons), which can exist separately of each other. In fact, in certain extreme environments such as neutron stars, extreme temperature and pressure prevents atoms from existing at all. The field of science which studies the true fundamental particles of matter is particle physics.

[edit] Philosophical atomism

Main article: Atomism

Up until the beginning of the 19th century, atomic theory was mainly philosophical and not founded in scientific experimentation. The earliest known theories were developed in ancient India in the 6th century BCE century BC by Hindu, Buddhist and Jaina philosophers. The first philosopher who formulated ideas about the atom in a systematic manner was Kanada. Another Indian philosopher, Pakudha Katyayana, who also lived in the 6th century BCE, had also propounded ideas about the atomic constitution of the material world. Indian atomists believed that an atom could be one of up to six elements, with each element having up to 24 properties. They developed detailed theories of how atoms could combine, react, vibrate, move, and perform other actions, and had particularly elaborate theories of how atoms combine, which explained how atoms first combine in pairs, and then group into trios of pairs, which are the smallest visible units of matter. They had also suggested the possibility of splitting an atom. (See Indian atomism for more details.)

Leucippus and Democritus, Greek philosophers in the 5th century BCE, presented their own theory of atoms. The Greeks believed that atoms were all made of the same material but had different shapes and sizes, which determined the physical properties of the material. For instance, the atoms of a liquid were thought to be smooth, allowing them to slide over each other.^[1] In this line of thought, graphite and diamond would be composed of two different kinds of atoms, but today we know that they're both isomers of carbon.

During the 11th century (in the Islamic Golden Age), Islamic atomists developed atomic theories that represent a synthesis of both Greek and Indian atomism. Older Greek and Indian ideas were further developed by Islamic atomists, along with new Islamic ideas, such as the possibility of there being particles smaller than an atom. As Islamic influence began spreading through Europe, the ideas of Islamic atomism, along with the older ideas of Greek and Indian atomism, spread throughout Europe by the end of the Middle Ages.

[edit] Modern atomic theory

[edit] Birth of modern atomic theory

Dalton used pictograms rather than letters to represent elements.

In the early years of the 19th century, John Dalton developed his atomic theory in which he proposed that each chemical element is composed of atoms of a single, unique type, and that though they are both immutable and indestructible, they can combine to form more complex structures (chemical compounds). How precisely Dalton arrived at his theory is not entirely clear, but nonetheless it allowed him to explain a number of chemistry puzzles he and his contemporaries were pondering at the time.

The first was the law of conservation of mass, formulated by Antoine Lavoisier in 1789, which states that the total mass in a chemical reaction remains constant (ie the reactants weigh the same as the products). This law suggested to Dalton that matter is fundamentally indestructible.

The second was the law of definite proportions. First proven by the French chemist Joseph Louis Proust in 1799,^[2] this law states that if a compound is broken down into its constituent elements, then the masses of the constituents will always have the same proportions, regardless of the quantity or source of the original substance. Proust had synthesized copper carbonate through numerous methods and found that in each case the ingredients combined in the same proportions as they were produced when he broke down natural copper carbonate.

Dalton studied and expanded upon Proust's work to develop the law of multiple proportions: if two elements form more than one compound between them, then the ratios of the masses of the second element which combine with a fixed mass of the first element will be ratios of small integers. One pair of reactions Dalton is believed to have studied involved nitric oxide (NO) and oxygen (O₂). In one combination, these gases formed dinitrogen trioxide (N₂O₃), but when he repeated the combination with twice the amount of oxygen (a ratio of 1:2 - small integers), they instead formed nitrogen dioxide (NO₂).

4NO + O₂ → 2N₂O₃

4NO + 2O₂ → 4NO₂

Another puzzle that Dalton pondered was why water absorbed different gases in different proportions; for example, he found that water absorbed carbon dioxide far better than it absorbed nitrogen. Dalton hypothesized this was due to the differences in mass and complexity of the gases' particles; indeed, carbon dioxide molecules (CO₂) are heavier and larger than nitrogen molecules (N₂).

In 1803 Dalton published his first list of relative atomic weights for a number of substances^[3] (though he did not publicly discuss how he obtained these figures until 1808). Dalton estimated the atomic weights according to the mass ratios in which they combined, with hydrogen being the basic unit. However, Dalton did not conceive that with some elements, atoms existed in molecules - eg pure oxygen exists as O₂. He also mistakenly believed that the simplest compound between any two elements is always one atom of each (so he thought water was HO, not H₂O). This, in addition to the crudity of his equipment, resulted in his table being highly flawed. For instance, he believed oxygen atoms were 5.5 times heavier than hydrogen atoms, because in water he measured 5.5 grams of oxygen for every 1 gram of hydrogen and believed the formula for water was HO (an oxygen atom is actually 16 times heavier than a hydrogen atom).

The flaw in Dalton's theory was corrected in 1811 by Amedeo Avogadro. Avogadro had 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 its volume).^[4] Avogadro's law allowed him to deduce the diatomic nature of numerous gases by studying the volumes at which they reacted. For instance: since two litres of hydrogen will react with just one litre of oxygen to produce two litres of water vapour (at constant pressure and temperature), it meant a single oxygen molecule splits in two in order to form two particles of water. Thus, Avogadro was able to offer more accurate estimates of the atomic mass of oxygen and various other elements, and firmly established the distinction between molecules and atoms.

In 1827, the British botanist Robert Brown observed that dust particles floating in water constantly jiggled about for no apparent reason. In 1905, Albert Einstein theorised that this Brownian motion was caused by the water molecules continuously knocking the grains about, and developed a hypothetical mathematical model to describe it.^[5] This model was validated experimentally in 1911 by French physicist Jean Perrin, thus providing additional validation for particle theory (and by extension atomic theory).

[edit] Discovery of subatomic particles

Thomson's cathode ray tube in which he observed the deflection of cathode rays by an electric field.

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 ray tubes.^[6] A cathode ray tube of the kind used by Thomson 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 an electric field (in addition to magnetic fields, which was already known). He concluded that these rays, rather than being waves, were composed of negatively charged particles he called "corpuscles" (they would later be renamed electrons by other scientists).

Thomson believed that the corpuscles emerged from the very atoms of the electrode. He thus concluded that atoms were divisible, and that the corpuscles were their building blocks. To explain the overall neutral charge of the atom, he proposed that the corpuscles were distributed in ring-like structures in a uniform sea or cloud of positive charge; this was the plum pudding model.^[7]

Since atoms were found to be actually divisible, physicists later invented the term "elementary particles" to describe for indivisible particles.

[edit] Discovery of the nucleus

The gold foil experiment
Top: Expected results: alpha particles passing through the plum pudding model of the atom undisturbed.
Bottom: Observed results: a small portion of the particles were deflected, indicating a small, concentrated positive charge.

Thomson's plum pudding model was disproved in 1909 by one of his 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.

In the gold foil experiment, Hans Geiger and Ernest Marsden shot alpha particles through a sheet of gold (striking a fluorescent screen that surrounded the foil).^[8] Given the very small mass of the electrons, the high mass and momentum of the alpha particles and the uniform distribution of positive charge of the plum pudding model, the experimenters expected all the alpha particles to either pass through without significant deflection or be absorbed. To their astonishment, a small fraction of the alpha particles experienced heavy deflection. This led Rutherford to propose the planetary model of the atom in which pointlike electrons orbited in the space around a massive, compact nucleus-like planets orbiting the Sun.^[9]

[edit] Discovery of isotopes

In 1913, Thomson channeled a stream of neon ions through magnetic and electric fields, striking a photographic plate on the other side. 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; thusly did he discover the existence of isotopes.^[10]

[edit] Discovery of nuclear particles

In 1918, Rutherford managed to split the nucleus of the atom when he bombarded nitrogen gas with alpha particles and observed hydrogen nuclei being emitted from the gas. Rutherford concluded that the hydrogen nuclei emerged from the nuclei of the nitrogen atoms themselves.^[11] He later found that the positive charge of any atom could always be equated to that of an integer number of hydrogen nuclei. This, coupled with the fact that hydrogen—being the lightest element ever found—had an atomic mass of 1, led him to conclude hydrogen nuclei were singular particles, a basic consituent of all atomic nuclei: the proton. 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 hitherto unknown neutrally charged particles.

In 1928, Walter Bothe observed that beryllium emitted an electrically neutral radiation when bombarded with alpha particles. In 1932, James Chadwick exposed various elements to this radiation, and by comparing the energies of recoiling charged particles from the different targets, he deduced that the radiation was composed of electrically neutral particles with a mass similar to that of a proton.^[12] Chadwick called these particles "neutrons".

[edit] Quantum physical models of the atom

The planetary model of the atom had shortcomings. Firstly, according to the Larmor formula in classical electromagnetism, an accelerating electric charge emits electromagnetic waves; an orbiting charge would steadily lose energy and spiral towards the nucleus, colliding with it in a tiny fraction of a second. Another phenomenon the model did not explain was why excited atoms only emit light with certain discrete spectra.

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 fixed amounts known as quanta (singular, quantum). In 1913, Niels Bohr incorporated this idea into his Bohr model of the atom, in which the electrons could only orbit the nucleus in particular circular orbits with fixed angular momentum and energy, their distances from the nucleus being proportional to their respective energies.^[13] Under this model atoms could not spiral into the nucleu because they could not lose energy in a continuous manner; instead, they could only make instantaneous "quantum leaps" between the fixed energy levels.^[14] 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).^[15] Bohr's model was extended by Arnold Sommerfeld in 1916 to include elliptical orbits, using a quantization of generalized momentum.

The ad hoc Bohr-Sommerfeld model was extremely difficult to use, but in return it made impressive predictions in agreement with certain spectral properties. However, the model was unable to explain multielectron atoms, predict transition rates or describe fine and hyperfine structure.

In 1924, Louis de Broglie proposed that all objects—particularly subatomic particles such as electrons—exhibit a degree of wave-like behavior. Erwin Schrodinger, fascinated by this idea, explored whether or not the movement of an electron in an atom could be better explained as a wave rather than as a particle. Schrodinger's equation, published in 1926^[16], describes an electron as a wavefunction instead of as a point particle, and it elegantly predicted many of the spectral phenomena Bohr's model failed to explain. Although this concept was mathematically convenient, it was difficult to visualize, and faced opposition.^[17] One of its critics, Max Born, proposed instead that Schrodinger's wavefunction described not the electron but rather all its possible states, and thus could be used to calculate the probability of finding an electron at any given location around the nucleus.^[18]

In 1927, Werner Heisenberg pointed out that since a wavefunction incorporates time as well as position, it is impossible to simultaneously derive precise values for both the position and momentum of a particle for any given point in time^[19]; this became known as the uncertainty principle.

The five atomic orbitals of a neon atom, separated and arranged in order of increasing energy. Each orbital holds up to two electrons, which exist for most of the time in the zones represented by the bubbles.

This new approach invaldiated 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—tends to exist more frequently in certain regions around the nucleus than others; this pattern is referred to as its atomic orbital.

[edit] See also

• Discoveries of the chemical elements
• History of thermodynamics
• Introduction to quantum mechanics
• Kinetic theory
• Quantum chemistry
• Timeline of microphysics
• Timeline of thermodynamics, statistical mechanics, and random processes

[edit] Notes

[edit] External links

• Ancient Atomism
• Discovery of the Neutron
• Atom: The Incredible World