In physics and chemistry, the atomic theory is a scientific theory of the nature of matter, which states that matter is composed of units called atoms. Atomic theory began as a philosophical concept in ancient Greece and entered the mainstream of the nineteenth century when discoveries in the field of chemistry showed that matter really behaves as if it were an atom.
The word atom originates in the atomic adjective of the ancient Greek, which means "indivisible". As explained in the history of nuclear energy. Nineteenth-century chemists began to use the term in relation to the growing number of irreducible chemical elements.
Apparently, in the early twentieth century, through various experiments with electromagnetism and radioactivity, physicists discovered that the so-called "indivisible atom" is actually a conglomeration of different subatomic particles (mainly electrons, protons and neutrons). These particles can exist separately. In fact, in extreme environments such as neutron stars, extreme temperature and pressure make the existence of atoms completely difficult.
Since atoms have been shown to be divisible, physicists subsequently invented the term "elementary particles" to describe the "indivisible", though not indestructible, parts of an atom. The scientific field that studies subatomic particles is particle physics, and in this field physicists hope to discover the true fundamental nature of matter.
The idea that matter is made up of discrete units is very old and occurs in many ancient cultures, such as Greece and India. However, these ideas were based on philosophical and theological reasoning, rather than on evidence and experimentation.
For this reason, they could not convince everyone, so atomism remained one of the many conflicting hypotheses about the nature of matter. It was only in the nineteenth century that the idea was accepted and refined by scientists, the science of chemistry, at first, producing discoveries that could be easily explained using the concept of atoms.
Near the end of the eighteenth century, two laws emerged on chemical reactions, without reference to the notion of atomic theory. 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 reaction products).
The second was the law of definite proportions. First demonstrated by the French chemist Joseph Louis Proust in 1799, this law states that, if a compound is broken down into its constituent elements, the masses of the constituent compounds will always have the same proportions, regardless of the quantity or source of the initial substance.
John Dalton studied and expanded the previous result and developed the law of multiple proportions: if two elements can be combined to form a series of compounds, the mass ratio of the second element that is combined with a fixed mass of the first element will be a relationship of small integers.
For example, Proust studied tin oxides and found that their masses were 88.1% tin and 11.9% oxygen, or 78.7% tin and 21.3% oxygen (these are tin oxides and tin dioxide), respectively. Dalton observed in these percentages that 100 g of tin will be combined with 13.5 g or 27 g of oxygen; 13.5 and 27 are in a ratio of 1: 2. Dalton discovered that an atomic theory of matter could elegantly explain this common pattern in chemistry. In the case of Proust's tin oxides, a tin atom will be combined with one or two oxygen atoms.
Dalton believed that the atomic theory could explain why water absorbs different gases in different proportions; For example, he discovered that water absorbs carbon dioxide much better than nitrogen. Dalton hypothesized that this is due to differences in mass and complexity of the particulate gas. In fact, the molecules of carbon dioxide (CO2) are heavier and larger than the nitrogen molecules (N2).
Dalton proposed that each chemical element is composed of atoms of a single type, and although they can not be modified or destroyed by chemical means, they can combine to form more complex structures (chemical compounds). This marked the first truly scientific theory of the atom, because Dalton arrived at his conclusions by experimenting and analyzing the results in an empirical way.
In 1803, Dalton presented the first list of relative atomic masses orally for several substances. This article was published in 1805, but he did not discuss exactly how he obtained these figures. The method was discovered for the first time in 1807 by his knowledge, Thomas Thomson, in the third edition of his manual, A system of chemistry. Finally, Dalton published a complete presentation in his own book, A new system of chemical philosophy, 1808 and 1810.
Dalton estimated the atomic masses according to the relations in which the masses are combined with the hydrogen atom taken as a unit. However, Dalton did not conceive that, in some elements, there are more similar atoms in the molecules; for example, pure oxygen exists as O2. Also, he mistakenly believes that the simplest compound between two elements is always one of each atom (therefore, it was thought that HO water, not H 2 O).
This, in addition to the rudeness of his team, has hindered his results. For example, in 1803 he believed that 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, therefore, he believed that the water formula was HO. By adopting better data, in 1806 he came to the conclusion that the atomic mass of oxygen should be 7, instead of 5.5, and he considered that weight for the rest of his life. Others, at that point, had already reached the conclusion that the oxygen atom should weigh 8 if the hydrogen weighs 1, assuming the formula of the water molecule (HO) of Dalton, or 16 assuming the modern water formula ( H2O).
The defect of Dalton's theory was corrected in principle in 1811 by Amedeo Avogadro. Avogadro proposed that equal volumes of gases, at equal temperature and pressure, contain the same number of molecules (that is, the mass of gas particles does not affect the volume it occupies). Avogadro's law allowed him to deduce the diatomic nature of many gases by studying the volumes in which they react. For example, when two liters of hydrogen react with only one liter of oxygen to produce two liters of water vapor (under constant pressure and temperature), it means that an oxygen molecule is divided in two to help form a two water particles. . Therefore, Avogadro was able to provide more accurate estimates of the atomic mass of oxygen and other elements, and established a clear distinction between molecules and atoms.
The Brownian movement
In 1827, the English botanist Robert Brown noticed that the dust particles inside the pollen granules that floated in the water were constantly stirring for no apparent reason. In 1905, Albert Einstein theorized that this Brownian motion was caused by constant clashes of pearls of water molecules, and developed a hypothetical mathematical model to describe the phenomenon. This model was validated experimentally in 1908 by the French physicist Jean Perrin, thus providing an additional validation of particle theory (and, by extension, atomic theory).
The discovery of subatomic particles
It was thought that atoms were the smallest possible division of matter until 1897, when JJ Thomson discovered the electron working with cathode rays.
A Crookes tube is a sealed glass vessel in which two electrodes are separated from the vacuum. When a potential difference is applied to the electrodes, cathode rays are generated, creating a bright area where they hit the glass at the opposite end of the tube. Experimentally, Thomson discovered that the rays could deviate from an electric field (in addition to the magnetic fields, which are already known). He came to the conclusion that these rays, instead of being a form of light, are actually composed of very little charged particles, which he called "corpuscles" (which would later be called "electrons" by other scientists). He measured the charge ratio of electrical mass and discovered that it was 1,800 times smaller than hydrogen, the smallest atom. These corpuscles were a particle completely different from those previously known.
Thomson has suggested that atoms are actually divisible, and that corpuscles are their constituent elements. To explain that the atom is an electrically neutral whole, it has advanced the hypothesis that the corpuscles are distributed in a great uniform of positive charges; this was the plum pudding model, in which the electrons were embedded in plums positively charged in a plum pudding (although in the Thomson model they were not stationary).
Discovering the nucleus
Thomson's model was infringed in 1909 by one of his former students, Ernest Rutherford, who discovered that most of the mass and positive charge of the atom are concentrated in a very small fraction of its volume in an area it assumed. which is in the center.
In the Geiger-Marsden experiment, Hans Geiger and Ernest Marsden (Rutherford's colleagues who worked on his suggestion) propelled alpha particles into thin metal sheets and measured their deformation using a fluorescence screen. Given the very small mass of the electrons, the high pulse of the alpha particles and the low concentration of the positive charge in the plum pudding model, the experimenters expected all the alpha particles to pass through the metal sheet without significant deviations . To his surprise, a small fraction of the alpha particles were strongly deviated. Rutherford concluded that the positive charge of the atom must be concentrated in a very small volume that produces an electric field strong enough to deflect the alpha particles so strongly.
This led Rutherford to propose a planetary model in which a cloud of electrons surrounds a small, compact nucleus of positive charge. Only such a concentration of charge could produce electric fields strong enough to cause large deviations.
The first steps towards a quantum model of the atom
The atomic planetary model had two significant deficiencies. The first was that, unlike the planets that orbit around a sun, electrons are charged particles. It is known that an electric charge of the accelerator emits electromagnetic waves according to the Larmor formula of classical electromagnetism. A task in orbit should constantly lose energy and spiral into the nucleus, colliding with it in a fraction of a second. The second problem was that the planetary model could not explain the emission and absorption spectra of the observed atoms.
Quantum theory revolutionized physics in the early twentieth century, when Max Planck and Albert Einstein postulated that light energy is emitted or absorbed in discrete quantities known as quanta. In 1913, Niels Bohr incorporated this idea into Bohr's atom model, the electrons can only revolve around the nucleus in some circular orbit with the angular momentum and the fixed distance energy from the nucleus (ie, range) is proportional to the energy. In this model, an electron can not enter the nucleus, because it could not lose energy in a continuous form; Instead, he could only instant "quantum leaps" between fixed energy levels. When this occurs, the light is emitted or absorbed at a frequency proportional to the energy difference (from where the absorption and emission of light in discrete spectra).
Bohr's model was not perfect. I could only predict the spectral lines of hydrogen; I could not predict those of atoms with more electrons. Worse still, as the spectrographic technology evolved, additional spectral lines of hydrogen were observed, lines that the Bohr model could not explain. In 1916, Arnold Sommerfeld added elliptical orbits to Bohr to explain the additional emission lines, but this made the model very difficult to use without being able to explain the more complex atoms.
The discovery of isotopes
When experimenting with radioactive decay products, in 1913, the radiochemist Frederick Soddy discovered that they appeared to be more than one element per position in the periodic table. The term isotope was invented by Margaret Todd as the proper name of these elements.
In the same year, JJ Thomson conducted an experiment in which he channeled a flow of neon ions through magnetic and electric fields, hitting a photographic plate at the other end. He noticed two bright spots on the plate, suggesting two different types of deflection trajectories. Thomson concluded that this is because some of the neon ions have another table. The nature of these different masses would be explained later by the discovery of neutrons in 1932.
Discovery of nuclear particles
In 1917 Rutherford bombarded nitrogen gas with alpha particles and found that the gas coming out of the hydrogen nuclei (Rutherford has recognized, as he previously obtained by bombarding the hydrogen atoms with alpha particles, and observing the nuclei of hydrogen in the products). Rutherford concluded that the hydrogen nuclei resulted from the nuclei of the nitrogen atoms (basically, it split the nitrogen atom).
Discovery of protons
Own work and that of their students, Bohr and Henry Moseley, Rutherford knew that the positive charge of any atom could always be equated with a hydrogen whole nucleus. This, coupled with the fact that the atomic mass of several elements is roughly equivalent with a number of hydrogen atoms - then assumed to be the easiest particles - have led to the conclusion that hydrogen nuclei are Singular particles and basic constituent of all the atomic nuclei. He called these proton particles.
Discovery of neutrons
Several experiments conducted by Rutherford have shown that the nuclear masses of most atoms surpass that of the protons they possess; he speculated that this mass surplus is composed of some unknown, electrically neutral particles, which provisionally he called "neutrons".
In 1928, Walter Bothe noticed that the beryl emits electrically neutral and very penetrating radiation when it is bombarded with alpha particles. It was later discovered that this radiation could remove the hydrogen atoms from the paraffin wax. Initially, it was thought to be high-energy gamma radiation, since gamma radiation had a similar effect on metallic electrons, but James Chadwick discovered that the ionizing effect is too strong to be caused by electromagnetic radiation, as long as the energy and momentum are preserved in interaction.
In 1932, Chadwick exposed several elements, such as hydrogen and nitrogen, to the mysterious "beryllium radiation", and when measuring the energies of charged particles, he deduced that the radiation actually consists of neutral electrical particles that could not be dough. like gamma rays, but it had to have a mass similar to that of a proton. Chadwick now claimed that these particles are Rutherford's neutrons. For the discovery of the neutron, Chadwick received the Nobel Prize in 1935.
Quantum model of the atom
In 1924, Louis de Broglie advanced the hypothesis that all moving particles, especially subatomic particles, such as electrons, exhibit some form of wave. Erwin Schrödinger, fascinated by this idea, explored whether the movement of an electron in an atom could be better explained as a wave than as a particle. The Schrödinger equation, published in 1926, describes an electron as a wave instead of a point particle. This approach has elegantly predicted many of the spectral phenomena that the Bohr model could not explain. Although this concept was mathematically convenient, it was difficult to visualize and faced opposition. One of his critics, Max Born, instead suggested that the Schrödinger wave function does not describe the electron, but all its possible states, and therefore could be used to calculate the probability of finding an electron. anywhere around the nucleus. This interpretation reconciled the two opposing theories of the nature of particles and waves, and introduced the idea of wave-particle duality. This theory states that the electron can exhibit both wavelength and particle properties. For example, it can be refracted as a wave, and it has mass as a particle.
A consequence of the description of electrons as a wave is the mathematical impossibility of simultaneously calculating the position and momentum of an electron. This was known as Heisenberg's uncertainty principle after the physicist Werner Heisenberg, who described and published it for the first time in 1927. He invalidated the Bohr model with its clear and 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 be found at any distance from the nucleus, but, depending on its energy level, it occurs more frequently in some regions around the nucleus than in others; this probability pattern is called atomic orbital.