ATOMIC THEORY

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 various new discoveries in chemistry that he and his contemporaries made.
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 (that is, the reactants have the same mass as the products).[7] 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,[8] 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 studied involved the combinations of "nitrous air", or what we now call nitric oxide (NO), and oxygen (O2). Under certain conditions, these gases formed an unknown product at a certain combining ratio (now known to be nitrogen dioxide (NO2)), but when he repeated the reaction under other conditions, exactly twice the amount of nitric oxide (a ratio of 1:2 -- small integers) reacted completely with oxygen to form a different product -- now known as dinitrogen trioxide (N2O3).
2NO + O2 → 2NO2
4NO + O2 → 2N2O3
Dalton also believed atomic theory could explain 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' respective particles. Indeed, carbon dioxide molecules (CO2) are heavier and larger than nitrogen molecules (N2).


Various atoms and molecules as depicted in John Dalton's A New System of Chemical Philosophy (1808).
In 1803 Dalton orally presented his first list of relative atomic weights for a number of substances. This paper was published in 1805, but he did not discuss there exactly how he obtained these figures.[9] The method was first revealed in 1807 by his acquaintance Thomas Thomson, in the third edition of Thomson's textbook, A System of Chemistry. Finally, Dalton published a full account in his own textbook, A New System of Chemical Philosophy, 1808 and 1810.
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 exist in molecules – e.g. 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).[10] 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).[11] 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 liters of hydrogen will react with just one liter of oxygen to produce two liters of water vapor (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 pollen particles 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 hypothetical mathematical model to describe it.[12] This model was validated experimentally in 1908 by French physicist Jean Perrin, thus providing additional validation for particle theory (and by extension atomic theory).
[edit] Discovery of subatomic particles


Thomson's Crookes tube in which he observed the deflection of cathode rays by an electric field. The purple line represents the deflected electron stream.
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.[13] 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 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 a uniform sea or cloud of positive charge; this was the plum pudding model.[14]
Since atoms were found to be actually divisible, physicists later invented the term "elementary particles" to describe 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 with negligible deflection.
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 (colleagues of Rutherford working at his behest) shot alpha particles through a thin sheet of gold, striking a fluorescent screen that surrounded the sheet.[15] Given the very small mass of the electrons, the high momentum of the alpha particles and the unconcentrated 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.[16]
[edit] Discovery of isotopes
While experimenting with the products of radioactive decay, in 1913 radiochemist Frederick Soddy discovered that there appeared to be more than one element at each position on the atomic table.[17] The term isotope was coined by Margaret Todd as a suitable name for these elements.
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.[18] The nature of this differing mass would later be explained by the discovery of neutrons in 1932.
[edit] Discovery of nuclear particles
See also: Prout's hypothesis
In 1918, Rutherford 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 (in effect, he split the atom).[19] 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 facts that hydrogen was the lightest element known and that the atomic mass of every other element was roughly equivalent to a whole multiple of hydrogen's atomic mass, led him to conclude hydrogen nuclei were singular particles and a basic constituent 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, 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 ionisation effect was too strong for it to be due to electromagnetic radiation. In 1932, he 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 with a mass similar to that of a proton.[20] For his discovery of the neutron, Chadwick received the Nobel Prize in 1935.
[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 small fraction of a second. Another phenomenon the model did not explain was why excited atoms only emit light within 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 discrete 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.[21] Under this model electrons could not spiral into the nucleus because they could not lose energy in a continuous manner; instead, they could only make instantaneous "quantum leaps" between the fixed energy levels.[21] 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).[21]
Bohr's model was only able to predict the spectral lines of hydrogen; it couldn't predict those of multielectron atoms. Worse still, as spectrographic technology improved, additional spectral lines in hydrogen were observed which 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 complex atoms.
In 1924, Louis de Broglie proposed that all moving particles–particularly subatomic particles such as electrons–exhibit a degree of wave-like behavior. Erwin Schrödinger, 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. Schrödinger's equation, published in 1926,[22] 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.[23] One of its critics, Max Born, proposed instead that Schrödinger'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.[24]

The five 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 exist for most of the time in the zones represented by the colored bubbles. Each electron is equally in both orbital zones, shown here by color only to highlight the different wave phase.
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; this became known as the uncertainty principle. This invalidated 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.

























Dalton's Atomic Theory
Democritus first suggested the existence of the atom but it took almost two millennia before the atom was placed on a solid foothold as a fundamental chemical object by John Dalton (1766-1844). Although two centuries old, Dalton's atomic theory remains valid in modern chemical thought.
• Elements are made of tiny particles called atoms.
• All atoms of a given element are identical.
• The atoms of a given element are different from those of any other element; the atoms of different elements can be distinguished from one another by their respective relative weights.
• Atoms of one element can combine with atoms of other elements to form chemical compounds; a given compound always has the same relative numbers of types of atoms.
• Atoms cannot be created, divided into smaller particles, nor destroyed in the chemical process; a chemical reaction simply changes the way atoms are grouped together.
Dalton's Atomic Theory
1) All matter is made of atoms. Atoms are indivisible and indestructible.
2) All atoms of a given element are identical in mass and properties
3) Compounds are formed by a combination of two or more different kinds of atoms.
4) A chemical reaction is a rearrangement of atoms.
Modern atomic theory is, of course, a little more involved than Dalton's theory but the essence of Dalton's theory remains valid. Today we know that atoms can be destroyed via nuclear reactions but not by chemical reactions. Also, there are different kinds of atoms (differing by their masses) within an element that are known as "isotopes", but isotopes of an element have the same chemical properties.
Many heretofore unexplained chemical phenomena were quickly explained by Dalton with his theory. Dalton's theory quickly became the theoretical foundation in chemistry.








Ernest Rutherford publishes his atomic theory describing the atom as having a central positive nucleus surrounded by negative orbiting electrons. This model suggested that most of the mass of the atom was contained in the small nucleus, and that the rest of the atom was mostly empty space. Rutherford came to this conclusion following the results of his famous gold foil experiment. This experiment involved the firing of radioactive particles through minutely thin metal foils (notably gold) and detecting them using screens coated with zinc sulfide (a scintillator). Rutherford found that although the vast majority of particles passed straight through the foil approximately 1 in 8000 were deflected leading him to his theory that most of the atom was made up of 'empty space'.

Ernest Rutherford publishes his atomic theory describing the atom as having a central positive nucleus surrounded by negative orbiting electrons. This model suggested that most of the mass of the atom was contained in the small nucleus, and that the rest of the atom was mostly empty space. Rutherford came to this conclusion following the results of his famous gold foil experiment. This experiment involved the firing of radioactive particles through minutely thin metal foils (notably gold) and detecting them using screens coated with zinc sulfide (a scintillator). Rutherford found that although the vast majority of particles passed straight through the foil approximately 1 in 8000 were deflected leading him to his theory that most of the atom was made up of 'empty space’