ATOMIC STRUCTURE

ATOMIC STRUCTURE

The atom is a basic unit of matter consisting of a dense, central nucleus surrounded by a cloud of negatively charged electrons. The atomic nucleus contains a mix of positively charged protons and electrically neutral neutrons (except in the case of Hydrogen-1, which is the only stable nuclide with no neutron). The electrons of an atom are bound to the nucleus by the electromagnetic force. Likewise, a group of atoms can remain bound to each other, forming a molecule. An atom containing an equal number of protons and electrons is electrically neutral, otherwise it has a positive or negative charge and is an ion. An atom is classified according to the number of protons and neutrons in its nucleus: the number of protons determines the chemical element, and the number of neutrons determine the isotope of the element.
The name atom comes from the Greek ἄτομος/átomos, α-τεμνω, which means uncuttable, something that cannot be divided further. The concept of an atom as an indivisible component of matter was first proposed by early Indian and Greek philosophers. In the 17th and 18th centuries, chemists provided a physical basis for this idea by showing that certain substances could not be further broken down by chemical methods. During the late 19th and early 20th centuries, physicists discovered subatomic components and structure inside the atom, thereby demonstrating that the atom was not indivisible. The principles of quantum mechanics were used to successfully model the atom.[1][2]
Relative to everyday experience, atoms are minuscule objects with proportionately tiny masses. Atoms can only be observed individually using special instruments such as the scanning tunneling microscope. Over 99.9% of an atom's mass is concentrated in the nucleus,[note 1] with protons and neutrons having roughly equal mass. Each element has at least one isotope with unstable nuclei that can undergo radioactive decay.

This can result in a transmutation that changes the number of protons or neutrons in a nucleus.[3] Electrons that are bound to atoms possess a set of stable energy levels, or orbitals, and can undergo transitions between them by absorbing or emitting photons that match the energy differences between the levels. The electrons determine the chemical properties of an element, and strongly influence an atom's magnetic properties.
History
Various atoms and molecules as depicted in John Dalton's A New System of Chemical Philosophy (1808). In 1803, English instructor and natural philosopher John Dalton used the concept of atoms to explain why elements always react in a ratio of small whole numbers—the law of multiple proportions—and why certain gases dissolve better in water than others.

Additional validation of particle theory (and by extension atomic theory) occurred in 1827 when botanist Robert Brown used a microscope to look at dust grains floating in water and discovered that they moved about erratically—a phenomenon that became known as "Brownian motion". J. Desaulx suggested in 1877 that the phenomenon was caused by the thermal motion of water molecules, and in 1905 Albert Einstein produced the first mathematical analysis of the motion.[11][12][13] French physicist Jean Perrin used Einstein's work to experimentally determine the mass and dimensions of atoms, thereby conclusively verifying Dalton's atomic theory.[14]
The physicist J. J. Thomson, through his work on cathode rays in 1897, discovered the electron and its subatomic nature, which destroyed the concept of atoms as being indivisible units.[15] Thomson believed that the electrons were distributed throughout the atom, with their charge balanced by the presence of a uniform sea of positive charge (the plum pudding model).
However, in 1909, researchers under the direction of physicist Ernest Rutherford bombarded a sheet of gold foil with helium ions and discovered that a small percentage were deflected through much larger angles than was predicted using Thomson's proposal. Rutherford interpreted the gold foil experiment as suggesting that the positive charge of an atom and most of its mass was concentrated in a nucleus at the center of the atom (the Rutherford model), with the electrons orbiting it like planets around a sun. Positively charged helium ions passing close to this dense nucleus would then be deflected away at much sharper angles.[16]

While experimenting with the products of radioactive decay, in 1913 radiochemist Frederick Soddy discovered that there appeared to be more than one type of atom at each position on the periodic table.[17] The term isotope was coined by Margaret Todd as a suitable name for different atoms that belong to the same element. J.J. Thomson created a technique for separating atom types through his work on ionized gases, which subsequently led to the discovery of stable isotopes.[18]
A Bohr model of the hydrogen atom, showing an electron jumping between fixed orbits and emitting a photon of energy with a specific frequency. Meanwhile, in 1913, physicist Niels Bohr revised Rutherford's model by suggesting that the electrons were confined into clearly defined, quantized orbits, and could jump between these, but could not freely spiral inward or outward in intermediate states.[19] An electron must absorb or emit specific amounts of energy to transition between these fixed orbits. When the light from a heated material was passed through a prism, it produced a multi-colored spectrum. The appearance of fixed lines in this spectrum was successfully explained by the orbital transitions.[20]
Chemical bonds between atoms were now explained, by Gilbert Newton Lewis in 1916, as the interactions between their constituent electrons.[21] As the chemical properties of the elements were known to largely repeat themselves according to the periodic law,[22] in 1919 the American chemist Irving Langmuir suggested that this could be explained if the electrons in an atom were connected or clustered in some manner. Groups of electrons were thought to occupy a set of electron shells about the nucleus.[23]
The Stern–Gerlach experiment of 1922 provided further evidence of the quantum nature of the atom. When a beam of silver atoms was passed through a specially-shaped magnetic field, the beam was split based on the direction of an atom's angular momentum, or spin. As this direction is random, the beam could be expected to spread into a line. Instead, the beam was split into two parts, depending on whether the atomic spin was oriented up or down.[24]
In 1926, Erwin Schrödinger, using Louis de Broglie's 1924 proposal that particles behave to an extent like waves, developed a mathematical model of the atom that described the electrons as three-dimensional waveforms, rather than point particles. A consequence of using waveforms to describe electrons is that it is mathematically impossible to obtain precise values for both the position and momentum of a particle at the same time; this became known as the uncertainty principle, formulated by Werner Heisenberg in 1926. In this concept, for each measurement of a position one could only obtain a range of probable values for momentum, and vice versa. Although this model was difficult to visualize, it was able to explain observations of atomic behavior that previous models could not, such as certain structural and spectral patterns of atoms larger than hydrogen. Thus, the planetary model of the atom was discarded in favor of one that described atomic orbital zones around the nucleus where a given electron is most likely to exist.[25][26
In the 1950s, the development of improved particle accelerators and particle detectors allowed scientists to study the impacts of atoms moving at high energies.[29] Neutrons and protons were found to be hadrons, or composites of smaller particles called quarks. Standard models of nuclear physics were developed that successfully explained the properties of the nucleus in terms of these sub-atomic particles and the forces that govern their interactions.[30]
Around 1985, Steven Chu and co-workers at Bell Labs developed a technique for lowering the temperatures of atoms using lasers. In the same year, a team led by William D. Phillips managed to contain atoms of sodium in a magnetic trap. The combination of these two techniques and a method based on the Doppler effect, developed by Claude Cohen-Tannoudji and his group, allows small numbers of atoms to be cooled to several microkelvin. This allows the atoms to be studied with great precision, and later led to the discovery of Bose-Einstein condensation.[31]

Components Subatomic particles
Nucleus
The binding energy needed for a nucleon to escape the nucleus, for various isotopes. All the bound protons and neutrons in an atom make up a tiny atomic nucleus, and are collectively called nucleons. The radius of a nucleus is approximately equal to fm, where A is the total number of nucleons.[39] This is much smaller than the radius of the atom, which is on the order of 105 fm.

The nucleons are bound together by a short-ranged attractive potential called the residual strong force. At distances smaller than 2.5 fm this force is much more powerful than the electrostatic force that causes positively charged protons to repel each other.[40]
For atoms with low atomic numbers, a nucleus that has a different number of protons than neutrons can potentially drop to a lower energy state through a radioactive decay that causes the number of protons and neutrons to more closely match. As a result, atoms with roughly matching numbers of protons and neutrons are more stable against decay. However, with increasing atomic number, the mutual repulsion of the protons requires an increasing proportion of neutrons to maintain the stability of the nucleus, which modifies this trend. Thus, there are no stable nuclei with equal proton and neutron numbers above atomic number Z = 20 (calcium); and as Z increases toward the heaviest nuclei, the ratio of neutrons per proton required for stability increases to about 1.5.[42] Illustration of a nuclear fusion process that forms a deuterium nucleus, consisting of a proton and a neutron, from two protons. A positron (e+)—an antimatter electron—is emitted along with an electron neutrino.
The number of protons and neutrons in the atomic nucleus can be modified, although this can require very high energies because of the strong force. Nuclear fusion occurs when multiple atomic particles join to form a heavier nucleus, such as through the energetic collision of two nuclei. For example, at the core of the Sun protons require energies of 3–10 keV to overcome their mutual repulsion—the coulomb barrier—and fuse together into a single nucleus.[43] Nuclear fission is the opposite process, causing a nucleus to split into two smaller nuclei—usually through radioactive decay. The nucleus can also be modified through bombardment by high energy subatomic particles or photons. If this modifies the number of protons in a nucleus, the atom changes to a different chemical element.[44][45] If the mass of the nucleus following a fusion reaction is less than the sum of the masses of the separate particles, then the difference between these two values is emitted as energy, as described by Albert Einstein's mass–energy equivalence formula, E = mc2, where m is the mass loss and c is the speed of light. This deficit is the binding energy of the nucleus.[46]
Electron cloud
Main articles: Electron cloud and Atomic orbital

A potential well, showing the minimum energy V(x) needed to reach each position x. A particle with energy E is constrained to a range of positions between x1 and x2.The electrons in an atom are attracted to the protons in the nucleus by the electromagnetic force. This force binds the electrons inside an electrostatic potential well surrounding the smaller nucleus, which means that an external source of energy is needed in order for the electron to escape.
Wave functions of the first five atomic orbitals. The three 2p orbitals each display a single angular node that has an orientation and a minimum at the center. Each atomic orbital corresponds to a particular energy level of the electron. The electron can change its state to a higher energy level by absorbing a photon with sufficient energy to boost it into the new quantum state. Likewise, through spontaneous emission, an electron in a higher energy state can drop to a lower energy state while radiating the excess energy as a photon. These characteristic energy values, defined by the differences in the energies of the quantum states, are responsible for atomic spectral lines.[49]
The amount of energy needed to remove or add an electron (the electron binding energy) is far less than the binding energy of nucleons. For example, it requires only 13.6 eV to strip a ground-state electron from a hydrogen atom,[51] compared to 2.23 Mev for splitting a deuterium nucleus.[52] Atoms are electrically neutral if they have an equal number of protons and electrons. Atoms that have either a deficit or a surplus of electrons are called ions. Electrons that are farthest from the nucleus may be transferred to other nearby atoms or shared between atoms. By this mechanism, atoms are able to bond into molecules and other types of chemical compounds like ionic and covalent network crystals

Struktur Atomik
Atom merupakam suatu satuan dasar berarti terdiri dari suatu rapat, nukleus pusat yang dikelilingi oleh suatu elektron-elektron yang bermuatan negatif. Inti-atom berisi suatu campuran dari proton-proton yang bermuatan positif dan neutron-neutron netral (kecuali di dalam kasus dari Hydrogen-1, yang adalah satu-satunya nuklida yang stabil tanpa adanya neutron). Elektron-elektron dari suatu atom terikat pada nukleus oleh gaya elektromagnetik
Demikian juga, suatu kelompok atom-atom tinggal satu sama lain, membentuk suatu molekul. Satu atom yang berisi satu nomor sama dari proton-proton dan elektron-elektron adalah secara elektris netral, jika tidak itu mempunyai suatu muatan positif atau muatan negatif yang disebut ion.
Fisikawan J. J.Thomson, yang bekerja dengan sinar katode dalam 1897, menemukan elektron yang sifatnya sub-atom. Thomson percaya bahwa muatan positif dari suatu atom dan kebanyakan dari massa nya dipusatkan di suatu nukleus di pusat dari atom dengan elektron-elektron mengorbitkan nya seperti planet-planet di sekitar suatu matahari yang dikenal denga model kue kismis.
Sementara itu, pada tahun 1913, fisikawan Niels Bohr meninjau kembali model Rutherford dengan mengusulkan bahwa elektron-elektron menempati garis edar seperti tampak pada gambar.
Banyaknya proton-proton dan neutron-neutron di dalam inti-atom itu dapat dimodifikasi, meski ini dapat memerlukan tenaga-tenaga sangat tinggi oleh karena gaya yang kuat. Paduan inti terjadi ketika gabungan zarah-zarah atomis yang ganda untuk membentuk suatu nukleus yang lebih berat, seperti melalui tubrukan yang giat dua nucleus.
Jika massa dari nukleus yang mengikuti suatu reaksi fusi adalah kurang dari jumlah dari zarah-zarah yang terpisah, lalu perbedaan antara dua nilai-nilai ini dipancarkan energi yang sama, seperti yang digambarkan oleh rumus kelembaman massa Albert Einstein, E =mc2, di mana m adalah massa dan c adalah kelajuan cahaya.
Elektron
Masing-masing edar atom berpasangan dengan arah energi tertentu dari elektron. Elektron itu dapat mengubah keadaan nya pada suatu arah energi yang lebih tinggi menarik suatu foton dengan tenaga yang cukup untuk menaikkan tegangan nya ke dalam keadaan kuantum yang baru.