Atom

• 1. Philosophical Origins And Early Theory

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  The word atom comes from the ancient Greek term atomos, meaning uncuttable. Ancient philosophers like Leucippus and his student Democritus proposed that all matter consists of indivisible units called atomos. This idea appeared in many cultures independently. In ancient India, Kanāda of the Vaiśeşika school proposed indivisible particles known as paramānu. Later Buddhist Abhidharma traditions also articulated atomistic views emphasizing momentary and causally dependent material phenomena. These early concepts differed fundamentally from modern scientific understanding. They were speculative ideas developed without experimental or quantitative frameworks. The Greeks and Indians both believed matter was composed of extremely small, indivisible units. Modern atomic theory is not based on these old philosophical concepts.

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• 2. Discovery Of Subatomic Particles

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  In 1897, J. J. Thomson discovered that cathode rays could be deflected by electric and magnetic fields. He measured these particles to be 1,700 times lighter than hydrogen. Thomson called these new particles corpuscles before they were renamed electrons. He showed that electrons were identical to particles given off by photoelectric and radioactive materials. Atoms were not indivisible as scientists had thought. The atom was composed of electrons whose negative charge was balanced out by some source of positive charge. Ions are atoms which have an excess or shortage of electrons. Between 1908 and 1913, Ernest Rutherford and his colleagues Hans Geiger and Ernest Marsden performed experiments bombarding thin foils of metal with alpha particles. They spotted a small number of alpha particles being deflected by angles greater than 90 degrees. This shouldn't have been possible according to the Thomson model. Rutherford proposed that the positive charge of the atom is concentrated in a tiny volume at the center. In 1920, he named this particle proton. James Chadwick discovered neutrons in 1932 after observing beryllium emitted highly penetrating radiation when bombarded with alpha particles.

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• 3. Evolution Of Atomic Models

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  Thomson imagined that the balance of electrostatic forces would distribute electrons throughout a sphere. His model became popularly known as the plum pudding model. It was unable to predict properties like emission spectra and valencies. A problem in classical mechanics is that an accelerating charged particle radiates electromagnetic radiation. An electron orbiting a central charge should spiral down into that nucleus as it loses speed. In 1913, Niels Bohr proposed a new model where electrons could only do so in a finite set of orbits. Electrons could jump between these orbits only in discrete changes of energy corresponding to absorption or radiation of a photon. This quantization explained why electron orbits are stable. Bohr's model could only predict the emission spectra of hydrogen. Werner Heisenberg published the first consistent mathematical formulation of quantum mechanics in 1925. Louis de Broglie had proposed that all particles behave like waves to some extent one year earlier. Erwin Schrödinger used this idea to develop the Schrödinger equation in 1926. The planetary model of the atom was discarded in favor of one describing atomic orbital zones around the nucleus.

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• 4. Quantum Mechanics And Orbital Theory

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  In 1927, Werner Heisenberg formulated the uncertainty principle stating it is mathematically impossible to obtain precise values for both position and momentum of a particle at a given point in time. For a given accuracy in measuring a position one could only obtain a range of probable values for momentum. Atomic orbitals describe regions inside an electrostatic potential well where each electron forms a type of three-dimensional standing wave. Only a discrete set of these orbitals exist around the nucleus. Orbitals can have one or more ring or node structures and differ from each other in size, shape and orientation. Each atomic orbital corresponds to a particular energy level of the electron. An electron can change its state to a higher energy level by absorbing a photon with sufficient energy. Through spontaneous emission, an electron in a higher energy state can drop to a lower energy state while radiating excess energy as a photon. These characteristic energy values are responsible for atomic spectral lines. It requires only 13.6 eV to strip a ground-state electron from a hydrogen atom compared to 2.23 million eV for splitting a deuterium nucleus.

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• 5. Nuclear Structure And Radioactivity

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  All bound protons and neutrons in an atom make up a tiny atomic nucleus called nucleons. The radius of a nucleus is approximately equal to 1.2 femtometres times the cube root of the total number of nucleons. This is much smaller than the radius of the atom which is on the order of 105 femtometres. Nucleons are bound together by a short-ranged attractive potential called the residual strong force. At distances smaller than 2.5 femtometres this force is much more powerful than the electrostatic force causing positively charged protons to repel each other. Every element has one or more isotopes that have unstable nuclei subject to radioactive decay. Alpha decay occurs when the nucleus emits an alpha particle consisting of two protons and two neutrons. Beta decay results from a transformation of a neutron into a proton or a proton into a neutron. Gamma decay results from a change in the energy level of the nucleus to a lower state emitting electromagnetic radiation. Each radioactive isotope has a characteristic half-life determined by the time needed for half of a sample to decay. About 339 nuclides occur naturally on Earth, of which 251 have not been observed to decay.

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• 6. Chemical Bonding And Periodic Trends

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  Valency is the combining power of an element determined by the number of bonds it can form to other atoms. The outermost electron shell of an atom in its uncombined state is known as the valence shell. Electrons in that shell are called valence electrons. The number of valence electrons determines bonding behavior with other atoms. A transfer of a single electron between atoms forms bonds like those occurring in sodium chloride. Many elements display multiple valences or tendencies to share differing numbers of electrons. Carbon and organic compounds take many forms of electron-sharing. Elements at the far right of the periodic table have their outer shell completely filled resulting in chemically inert noble gases. Atoms lack a well-defined outer boundary so dimensions are described in terms of atomic radius. Atomic radii vary with location on the atomic chart and type of chemical bond. Helium has a radius of 32 pm while caesium reaches 225 pm. On the periodic table, atom size tends to increase when moving down columns but decrease when moving across rows.

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• 7. Cosmic Nucleosynthesis And Abundance

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  Electrons exist in the Universe since early stages of the Big Bang. In about three minutes Big Bang nucleosynthesis produced most helium, lithium, and deuterium. Atoms became to dominate over charged particles 380,000 years after the Big Bang during recombination. Since the Big Bang, atomic nuclei have been combined in stars through nuclear fusion to produce more helium and elements from carbon up to iron. Isotopes such as lithium-6 are generated in space through cosmic ray spallation. Elements heavier than iron were produced in supernovae and colliding neutron stars through the r-process. Elements such as lead formed largely through radioactive decay of heavier elements. Baryonic matter forms about 4% of total energy density of observable universe. Up to 95% of Milky Way baryonic matter is concentrated inside stars where conditions are unfavorable for atomic matter. The Earth contains approximately 10^50 atoms. Most helium in crust of Earth is product of alpha decay. Carbon-14 is continuously generated by cosmic rays in atmosphere.

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