Atom Definition and Atomic Theory: What is an Atom? Atom Structure and Atom Facts
The smallest component of an element having the chemical properties of the element, consisting of a nucleus containing combinations of neutrons and protons and one or more electrons bound to the nucleus by electrical attraction; the number of protons determines the identity of the element.
An atom with one of the electrons replaced by some other particle: muonic atom; kaonic atom.
Energy. This component as the source of nuclear energy.
A hypothetical particle of matter so minute as to admit of no division.
Anything extremely small; a minute quantity.
The main three particles that make up the atom are, in the nucleus, positively charged protons and neutral neutrons and, in the various electron shells surrounding the nucleus, negatively charged electrons.
The number of protons determines the atomic number of the atom and its properties. The number of neutrons can vary for any given atomic number, providing isotopes of the element. Electrons can be gained or lost to form negative or positive ions.
Protons, neutrons, and electrons are the basic components.
The atom is a basic unit of matter that consists 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 neutrons). 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 by chemical bonds based on the same force, forming a molecule. An atom containing an equal number of protons and electrons is electrically neutral, otherwise it is positively or negatively charged and is known as 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 determines the isotope of the element.
Chemical atoms, which in science now carry the simple name of "atom," are minuscule objects with diameters of a few tenths of a nanometer and tiny masses proportional to the volume implied by these dimensions. Atoms can only be observed individually using special instruments such as the scanning tunneling microscope. Over 99.94% of an atom's mass is concentrated in the nucleus, with protons and neutrons having roughly equal mass. Each element has at least one isotope with an unstable nucleus that can undergo radioactive decay. This can result in a transmutation that changes the number of protons or neutrons in a nucleus. 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. The principles of quantum mechanics have been successfully used to model the observed properties of the atom.
What are Atoms?
Atoms are the basic building blocks of matter that make up everyday objects. A desk, the air, even you are made up of atoms!
There are 90 naturally occurring kinds of atoms. Scientists in labs have been able to make about 25 more.
Atomic Theory and History
Until the final years of the nineteenth century, the accepted model of the atom resembled that of a billiard ball - a small, solid sphere. In 1897, J. J. Thomson dramatically changed the modern view of the atom with his discovery of the electron. Thomson's work suggested that the atom was not an "indivisible" particle as John Dalton had suggested but, a jigsaw puzzle made of smaller pieces.
Thomson's notion of the electron came from his work with a nineteenth century scientific curiosity: the cathode ray tube. For years scientists had known that if an electric current was passed through a vacuum tube, a stream of glowing material could be seen; however, no one could explain why. Thomson found that the mysterious glowing stream would bend toward a positively charged electric plate. Thomson theorized, and was later proven correct, that the stream was in fact made up of small particles, pieces of atoms that carried a negative charge. These particles were later named electrons.
After Eugen Goldstein’s 1886 discovery that atoms had positive charges, Thomson imagined that atoms looked like pieces of raisin bread, a structure in which clumps of small, negatively charged electrons (the "raisins") were scattered inside a smear of positive charges. In 1908, Ernest Rutherford, a former student of Thomson's, proved Thomson's raisin bread structure incorrect.
Rutherford performed a series of experiments with radioactive alpha particles. While it was unclear at the time what the alpha particle was, it was known to be very tiny. Rutherford fired tiny alpha particles at solid objects such as gold foil. He found that while most of the alpha particles passed right through the gold foil, a small number of alpha particles passed through at an angle (as if they had bumped up against something) and some bounced straight back like a tennis ball hitting a wall. Rutherford's experiments suggested that gold foil, and matter in general, had holes in it! These holes allowed most of the alpha particles to pass directly through, while a small number ricocheted off or bounced straight back because they hit a solid object.
In 1911, Rutherford proposed a revolutionary view of the atom. He suggested that the atom consisted of a small, dense core of positively charged particles in the center (or nucleus) of the atom, surrounded by a swirling ring of electrons. The nucleus was so dense that the alpha particles would bounce off of it, but the electrons were so tiny, and spread out at such great distances, that the alpha particles would pass right through this area of the atom. Rutherford's atom resembled a tiny solar system with the positively charged nucleus always at the center and the electrons revolving around the nucleus.
Figure 1: Interpreting Rutherford's Gold Foil Experiment
The positively charged particles in the nucleus of the atom were called protons. Protons carry an equal, but opposite, charge to electrons, but protons are much larger and heavier than electrons.
In 1932, James Chadwick discovered a third type of subatomic particle, which he named the neutron. Neutrons help stabilize the protons in the atom's nucleus. Because the nucleus is so tightly packed together, the positively charged protons would tend to repel each other normally. Neutrons help to reduce the repulsion between protons and stabilize the atom's nucleus. Neutrons always reside in the nucleus of atoms and they are about the same size as protons. However, neutrons do not have any electrical charge; they are electrically neutral.
Atoms are electrically neutral because the number of protons (+ charges) is equal to the number of electrons (- charges) and thus the two cancel out. As the atom gets larger, the number of protons increases, and so does the number of electrons (in the neutral state of the atom). The illustration linked below compares the two simplest atoms, hydrogen and helium.
Atoms are extremely small. One hydrogen atom (the smallest atom known) is approximately 5 x 10-8 mm in diameter. To put that in perspective, it would take almost 20 million hydrogen atoms to make a line as long as this dash -. Most of the space taken up by an atom is actually empty because the electron spins at a very far distance from the nucleus. For example, if we were to draw a hydrogen atom to scale and used a 1-cm proton (about the size of this picture - ), the atom's electron would spin at a distance of ~0.5 km from the nucleus. In other words, the atom would be larger than a football field!
Atoms of different elements are distinguished from each other by their number of protons (the number of protons is constant for all atoms of a single element; the number of neutrons and electrons can vary under some circumstances). To identify this important characteristic of atoms, the term atomic number (z) is used to describe the number of protons in an atom. For example, z = 1 for hydrogen and z = 2 for helium.
Another important characteristic of an atom is its weight, or atomic mass. The weight of an atom is roughly determined by the total number of protons and neutrons in the atom. While protons and neutrons are about the same size, the electron is more that 1,800 times smaller than the two. Thus the electrons' weight is inconsequential in determining the weight of an atom - it's like comparing the weight of a flea to the weight of an elephant. Refer to the animation above to see how the number of protons plus neutrons in the hydrogen and helium atoms corresponds to the atomic mass.
Atomic Theory Controversy
Dalton's atomic theory remained controversial throughout the 19th Century. Whilst the law of definite proportion were accepted, the hypothesis that this was due to atoms was not so widely accepted. For example in 1826 when Sir Humphry Davy presented Dalton the Royal Medal from the Royal Society, Davy said that the theory only became useful when the atomic conjecture was ignored.
Sir Benjamin Collins Brodie in 1866 published the first part of his Calculus of Chemical Operations as a non atomic alternative to the atomic theory. He described atomic theory as a 'Thoroughly materialistic bit of joiners work'. Alexander Williamson used his Presidential address to the London Chemical Society in 1869 to defend the atomic theory against its critics and doubters. This in turn led to further meetings at which positivists again attacked the supposition that there were atoms. Ernst Mach also opposed the atomic theory. The matter was finally resolved in Dalton's favor by Einstein's work on brownian motion in the early 20th Century.
In 1905, Albert Einstein demonstrated the physical reality of the photons, hypothesized by Max Planck in 1900, in order to solve the problem of black body radiation in thermodynamics.
In 1874, G. Johnstone Stoney postulated a minimum unit of electrical charge, for which he suggested the name electron in 1891. In 1897, J. J. Thomson confirmed Stoney's conjecture by discovering the first subatomic particle, the electron (now denoted e−). Subsequent speculation about the structure of atoms was severely constrained by Ernest Rutherford's 1907 gold foil experiment, showing that the atom is mainly empty space, with almost all its mass concentrated in a (relatively) tiny atomic nucleus. The development of the quantum theory led to the understanding of chemistry in terms of the arrangement of electrons in the mostly empty volume of atoms. In 1918, Rutherford confirmed that the hydrogen nucleus was a particle with a positive charge, which he named the proton, now denoted p+. Rutherford also conjectured that all nuclei other than hydrogen contain chargeless particles, which he named the neutron. It is now denoted n. James Chadwick discovered the neutron in 1932. The word nucleon denotes neutrons and protons collectively.
Neutrinos were postulated in 1931 by Wolfgang Pauli (and named by Enrico Fermi) to be produced in beta decays of neutrons, but were not discovered until 1956. Pions were postulated by Hideki Yukawa as mediators of the residual strong force, which binds the nucleus together. The muon was discovered in 1936 by Carl D. Anderson, and initially mistaken for the pion. In the 1950s the first kaons were discovered in cosmic rays.
The development of new particle accelerators and particle detectors in the 1950s led to the discovery of a huge variety of hadrons, prompting Wolfgang Pauli's remark: "Had I foreseen this, I would have gone into botany". The classification of hadrons through the quark model in 1961 was the beginning of the golden age of modern particle physics, which culminated in the completion of the unified theory called the standard model in the 1970s. The discovery of the weak gauge bosons through the 1980s, and the verification of their properties through the 1990s is considered to be an age of consolidation in particle physics. As of early 2012, of all the particles in the Standard Model, only the existence of the Higgs boson remained to be verified. On July 4th, 2012, CERN announced the discovery of a new particle, compatible with the Standard Model Higgs boson, through experiments conducted with the Large Hadron Collider.
Discovery of subatomic particles
In physics or chemistry, subatomic particles are the particles smaller than an atom.
In particle physics, the conceptual idea of a particle is one of several concepts inherited from classical physics. This describes the world we experience, used (for example) to describe how matter and energy behave at the molecular scales of quantum mechanics. For physicists, the word "particle" means something rather different from the common sense of the term, reflecting the modern understanding of how particles behave at the quantum scale in ways that differ radically from what everyday experience would lead us to expect.
The idea of a particle underwent serious rethinking in light of experiments that showed that light could behave like a stream of particles (called photons) as well as exhibit wave-like properties. These results necessitated the new concept of wave-particle duality to reflect that quantum-scale "particles" are understood to behave in a way resembling both particles and waves. Another new concept, the uncertainty principle, concluded that analyzing particles at these scales would require a statistical approach. In more recent times, wave-particle duality has been shown to apply not only to photons but to increasingly massive particles.
All of these factors ultimately combined to replace the notion of discrete "particles" with the concept of "wave-packets" of uncertain boundaries, whose properties are known only as probabilities, and whose interactions with other "particles" remain largely a mystery, even 80 years after the establishment of quantum mechanics.
What are the components of an atom? How much does each atom weigh?
All atoms are made up of three subatomic particles, the proton, neutron and electron
The proton has a +1 charge. It is a fairly heavy particle and resides in the dense nucleus of an atom
The electron has a -1 charge. It is a light particle with a mass ~1/2000 of a proton. It resides in a cloud around the nucleus
The neutron has a 0 charge. It is a heavy particle with a mass about equal to a proton and resides in the nucleus
Around 140 years ago Dimitri Ivanovich Mendeleyev put the names and properties of all the chemicals he knew onto cards. He then tried arranging them in a way that made sense in a sort of Chemist's game of Solitaire. The resulting chart, called The Periodic Table of Elements, has contributed probably more than anything else to our understanding of matter. When you begin studying the Periodic table in school, pay close attention. There is a tremendous amount of information and beauty tucked away on that single sheet if you understand the subtleties.
Ok, how does that answer your questions? The first question is easy. Atoms being made up of combinations of protons, electrons and usually, but not always, some neutrons thrown in for fun.
The answer to the second question is, "It depends on what atom it is." The weight of each element is different depending on the number of protons and neutrons it has. The elements in Mr. Mendeleyev's list are arranged by their atomic number, which is the number of protons in that element. Also listed is that elements atomic weight, which is really the average number of protons and neutrons. I hope you noticed I said average weight, because elements often exist in different varieties called isotopes. This means that they have a different number of neutrons. The atomic weight listed is the average of the atomic weights of the element and its isotopes. If the number of protons was different it would be a different element. The mass of the electrons is very small and ignored in all this. The number of electrons is the same as the number of protons, but all sorts of things in nature can change that. The unit of atomic weight was fixed at 1/12 the atomic weight of a carbon atom.
Now it gets tricky. The atomic weight times a constant called Avogadro's number (6.0225*1023) is the amount of atoms that have a mass in grams equivalent to the atomic weight. Let's pick an example of the third most common element in the universe, carbon, with an atomic weight of 12.01. That means 12.01 grams of carbon contains 6.0225*1023 atoms of carbon. If you divide Avogadro's number by the atomic weight you get the mass of one atom. In the case of carbon, that ends up being 0.00000000000000000000019942 grams. Try weighing that on a bathroom scale!
History of the periodic table
First systemization attempts
The discovery of the elements mapped to significant periodic table development dates (pre-, per- and post-)
In 1789, Antoine Lavoisier published a list of 33 chemical elements, grouping them into gases, metals, nonmetals, and earths; Chemists spent the following century searching for a more precise classification scheme. In 1829, Johann Wolfgang Döbereiner observed that many of the elements could be grouped into triads based on their chemical properties. Lithium, sodium, and potassium, for example, were grouped together in a triad as soft, reactive metals. Döbereiner also observed that, when arranged by atomic weight, the second member of each triad was roughly the average of the first and the third; this became known as the Law of Triads. German chemist Leopold Gmelin worked with this system, and by 1843 he had identified ten triads, three groups of four, and one group of five. Jean-Baptiste Dumas published work in 1857 describing relationships between various groups of metals. Although various chemists were able to identify relationships between small groups of elements, they had yet to build one scheme that encompassed them all.
In 1858, German chemist August Kekulé observed that carbon often has four other atoms bonded to it. Methane, for example, has one carbon atom and four hydrogen atoms. This concept eventually became known as valency; different elements bond with different numbers of atoms.
In 1862, Alexandre-Emile Béguyer de Chancourtois, a French geologist, published an early form of periodic table, which he called the telluric helix or screw. He was the first person to notice the periodicity of the elements. With the elements arranged in a spiral on a cylinder by order of increasing atomic weight, de Chancourtois showed that elements with similar properties seemed to occur at regular intervals. His chart included some ions and compounds in addition to elements. His paper also used geological rather than chemical terms and did not include a diagram; as a result, it received little attention until the work of Dmitri Mendeleev.
In 1864, Julius Lothar Meyer, a German chemist, published a table with 44 elements arranged by valency. The table showed that elements with similar properties often shared the same valency. Concurrently, William Odling (an English chemist) published an arrangement of 57 elements, ordered on the basis of their atomic weights. With some irregularities and gaps, he noticed what appeared to be a periodicity of atomic weights amongst the elements and that this accorded with 'their usually received groupings.' Odling alluded to the idea of a periodic law but did not pursue it. He subsequently proposed (in 1870) a valence-based classification of the elements.
English chemist John Newlands produced a series of papers in 1864 and 1865 noting that when the elements were listed in order of increasing atomic weight, similar physical and chemical properties recurred at intervals of eight; he likened such periodicity to the octaves of music. This Law of Octaves, however, was ridiculed by Newlands' contemporaries, and the Chemical Society refused to publish his work. Newlands was nonetheless able to draft a table of the elements and used it to predict the existence of missing elements, such as germanium. The Chemical Society only acknowledged the significance of his discoveries five years after they credited Mendeleev.
In 1867, Gustavus Hinrichs, a Danish born academic chemist based in America, published a spiral periodic system based on atomic spectra and weights, and chemical similarities. His work was regarded as idiosyncratic, ostentatious and labyrinthine and this may have militated against its recognition and acceptance.
Mendeleev's 1869 periodic table; note that his arrangement presents the periods vertically, and the groups horizontally.
Russian chemistry professor Dmitri Mendeleev and German chemist Julius Lothar Meyer independently published their periodic tables in 1869 and 1870, respectively. Mendeleev's table was his first published version; that of Meyer was an expanded version of his (Meyer's) table of 1864. They both constructed their tables by listing the elements in rows or columns in order of atomic weight and starting a new row or column when the characteristics of the elements began to repeat.
The recognition and acceptance afforded Mendeleev's table came from two decisions he made. The first was to leave gaps in the table when it seemed that the corresponding element had not yet been discovered. Mendeleev was not the first chemist to do so, but he was the first to be recognized as using the trends in his periodic table to predict the properties of those missing elements, such as gallium and germanium. The second decision was to occasionally ignore the order suggested by the atomic weights and switch adjacent elements, such as tellurium and iodine, to better classify them into chemical families. With the development of theories of atomic structure, it became apparent that Mendeleev had unintentionally listed the elements in order of increasing atomic number or nuclear charge.
The significance of atomic numbers to the organization of the periodic table was not appreciated until the existence and properties of protons and neutrons became understood. Mendeleev's periodic tables used atomic weight instead of atomic number to organize the elements, information determinable to fair precision in his time. Atomic weight worked well enough in most cases to (as noted) give a presentation that was able to predict the properties of missing elements more accurately than any other method then known. Substitution of atomic numbers, once understood, gave a definitive, integer-based sequence for the elements, still used today even as new synthetic elements are being produced and studied.
Short form of periodic table, as originally published by Mendeleev in 1871, updated with all elements discovered to 2012.
In 1871, Mendeleev published an updated form of periodic table (shown above), as well as giving detailed predictions for the elements he had earlier noted were missing, but should exist. These gaps were subsequently filled as chemists discovered additional naturally occurring elements. It is often stated that the last naturally occurring element to be discovered was francium (referred to by Mendeleev as eka-caesium) in 1939. However, plutonium, produced synthetically in 1940, was identified in trace quantities as a naturally occurring primordial element in 1971, and in 2011 it was found that all the elements up to californium can occur naturally as trace amounts in uranium ores by neutron capture and beta decay.
The popular periodic table layout, also known as the common or standard form, is attributable to Horace Groves Deming. In 1923, Deming, an American chemist, published short (Mendeleev style) and medium (18-column) form periodic tables. Merck and Company prepared a handout form of Deming's 18-column medium table, in 1928, which was widely circulated in American schools. By the 1930s Deming's table was appearing in handbooks and encyclopaedias of chemistry. It was also distributed for many years by the Sargent-Welch Scientific Company.
With the development of modern quantum mechanical theories of electron configurations within atoms, it became apparent that each period (row) in the table corresponded to the filling of a quantum shell of electrons. Larger atoms have more electron sub-shells, so later tables have required progressively longer periods.
Glenn T. Seaborg who, in 1945, suggested a new periodic table showing the actinides as belonging to a second f-block series
In 1945, Glenn Seaborg, an American scientist, made the suggestion that the actinide elements, like the lanthanides were filling an f sub-level. Before this time the actinides were thought to be forming a fourth d-block row. Seaborg's colleagues advised him not to publish such a radical suggestion as it would most likely ruin his career. As Seaborg considered he did not then have a career to bring into disrepute, he published anyway. Seaborg's suggestion was found to be correct and he subsequently went on to win the 1951 Nobel prize in chemistry for his work in synthesizing actinide elements.
Although minute quantities of some transuranic elements occur naturally, they were all first discovered in laboratories. Their production has expanded the periodic table significantly, the first of these being neptunium, synthesized in 1939. Because many of the transuranic elements are highly unstable and decay quickly, they are challenging to detect and characterize when produced. There have been controversies concerning the acceptance of competing discovery claims for some elements, requiring independent review to determine which party has priority, and hence naming rights. The most recently accepted and named elements are flerovium (element 114) and livermorium (element 116), both named on 31 May 2012. In 2010, a joint Russia–US collaboration at Dubna, Moscow Oblast, Russia, claimed to have synthesized six atoms of ununseptium (element 117), making it the most recently claimed discovery.