Atom
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"Atom" and "Atomic" redirect here. For other uses, see Atom (disambiguation).
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In chemistry and physics, an atom (Greek άτομον meaning "indivisible") is the smallest possible particle of a chemical element that retains its chemical properties. The word atom may also refer to the smallest possible indivisible fundamental particle. This definition must not be confused with that of chemical atoms, since chemical atoms (hereafter "atoms") are composed of smaller subatomic particles.
Most atoms are composed of three types of massive subatomic particles which govern their external properties:
- electrons, which have a negative charge and are the least massive of the three;
- protons, which have a positive charge and are about 1836 times more massive than electrons; and
- neutrons, which have no charge and are about 1838 times more massive than electrons.
Protons and neutrons are both nucleons and make up the dense, massive atomic nucleus. The electrons form the much larger electron cloud surrounding the nucleus.
Atoms differ in the number of each of the subatomic particles they contain. The number of protons in an atom (called the atomic number) determines the element of the atom. Within a single element, the number of neutrons may also vary, determining the isotope of that element. Atoms are electrically neutral if they have an equal number of protons and electrons. Electrons that are furthest from the nucleus may be transferred to other nearby atoms or even shared between atoms. Atoms which have either a deficit or a surplus of electrons are called ions. The number of protons and neutrons in the atomic nucleus may also change, via nuclear fusion, nuclear fission or radioactive decay.
Atoms are the fundamental building blocks of chemistry, and are conserved in chemical reactions. Atoms are able to bond into molecules and other types of chemical compounds. Molecules are made up of multiple atoms; for example, a molecule of water is a combination of two hydrogen atoms and one oxygen atom.
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Properties of the atom
Subatomic particles
- see main article subatomic particles
Although the name "atom" was applied at a time when atoms were thought to be indivisible, it is now known that the atom can be broken down into a number of smaller components. The first of these to be discovered was the negatively charged electron, which is easily ejected from atoms during ionization. The electrons orbit a small, dense body containing all of the positive charge in the atom, called the atomic nucleus. This nucleus is itself made up of nucleons: positively charged protons and chargeless neutrons.
Before 1961, the subatomic particles were thought to consist of only protons, neutrons and electrons. However, protons and neutrons themselves are now known to consist of still smaller particles called quarks. In addition, the electron is known to have a nearly massless neutral partner called a neutrino. Together, the electron and neutrino are both leptons.
Ordinary atoms are composed only of quarks and leptons of the first generation. The proton is composed of two up quarks and one down quark, whereas the neutron is composed of one up quark and two down quarks. Although they do not occur in ordinary matter, two other heavier generations of quarks and leptons may be generated in high-energy collisions.
The subatomic force carrying particles (called gauge bosons) are also important to atoms. Electrons are bound to the nucleus by photons carrying the electromagnetic force. Protons and neutrons are bound together in the nucleus by gluons carrying the strong nuclear force.
Electron configuration
- see main article electron configuration
The chemical behavior of atoms is due to interactions between electrons. Electrons of an atom remain within certain, predictable electron configurations. These configurations are determined by the quantum mechanics of electrons in the electric potential of the atom; the principal quantum number determines particular electron shells with distinct energy levels. Generally, the higher the energy level of a shell, the further away it is from the nucleus. The electrons in the outermost shell, called the valence electrons, have the greatest influence on chemical behavior. Core electrons (those not in the outer shell) play a role, but it is usually in terms of a secondary effect due to screening of the positive charge in the atomic nucleus.
An electron shell can hold up to 2n2 electrons, where n is the principal quantum number of the shell. The occupied shell of greatest n is the valence shell, even if it only has one electron. In the most stable ground state, an atom's electrons will fill up its shells in order of increasing energy. Under some circumstances an electron may be excited to a higher energy level (that is, it absorbs energy from an external source and leaps to a higher shell), leaving a space in a lower shell. An excited atom's electrons will spontaneously fall into lower levels, emitting excess energy as a photons, until it returns to the ground state.
In addition to its principal quantum number n, an electron is distinguished by three other quantum numbers: the azimuthal quantum number l (describing the orbital angular momentum of the electron), the magnetic quantum number m (describing the direction of the angular momentum vector), and the spin quantum number s (describing the direction of the electron's intrinsic angular momentum). Electrons with varying l and m have distinctive shapes denoted by spectroscopic notation. In the illustration, the letters s, p, d and f (corresponding to l = 0, 1, 2, 3) describe the shape of the atomic orbital. In most atoms, orbitals of differing l are not exactly degenerate but separated into a fine structure. Orbitals of differing m are degenerate but may be separated by applying a magnetic field, creating the Zeeman effect. Electrons with differing s have very slight energy differences called hyperfine splitting.
Nucleon properties
The constituent protons and neutrons of the atomic nucleus are collectively called nucleons. The nucleons are held together in the nucleus by the strong nuclear force.
Nuclei can undergo transformations that affect the number of protons and neutrons they contain, a process called radioactive decay. When nuclei transformations take place spontaneously, this process is called radioactivity. Radioactive transformations proceed by a wide variety of modes, but the most common are alpha decay (emission of a helium nucleus) and beta decay (emission of an electron). Decays involving electrons or positrons are due to the weak nuclear interaction.
In addition, like the electrons of the atom, the nucleons of nuclei may be pushed into excited states of higher energy. However, these transitions typically require thousands of times more energy than electron excitations. When an excited nucleus emits a photon to return to the ground state, the photon has very high energy and is called a gamma ray.
Nuclear transformations also take place in nuclear reactions. In nuclear fusion, two light nuclei come together and merge into a single heavier nucleus. In nuclear fission, a single large nucleus is divided into two or more smaller nuclei.
Atom size and speed
Atoms are much smaller than the wavelengths of light that human vision can detect, so atoms cannot be seen in any kind of optical microscope. However, there are ways of detecting the positions of atoms on the surface of a solid or a thin film so as to obtain images. These include: electron microscopes (such as in scanning tunneling microscopy (STM)), atomic force microscopy (AFM), nuclear magnetic resonance (NMR) and x-ray microscopy.
Since the electron cloud does not have a sharp cutoff, the size of an atom is not easily defined. For atoms that can form solid crystal lattices, the distance between the centers of adjacent atoms can be easily determined by x-ray diffraction, giving an estimate of the atoms' size. For any atom, one might use the radius at which the electrons of the valence shell are most likely to be found. As an example, the size of a hydrogen atom is estimated to be approximately 1.0586×10−10 m (twice the Bohr radius). Compare this to the size of the proton (the only particle in the nucleus of the hydrogen atom), which is approximately 10−15 m. So the ratio of the size of the hydrogen atom to its nucleus is about 100,000:1. If an atom were the size of a stadium, the nucleus would be the size of a marble. Nearly all the mass of an atom is in its nucleus, yet almost all the space in an atom is filled by its electrons.
Atoms of different elements do vary in size, but the sizes do not scale linearly with the mass of the atom. Their sizes are roughly the same to within a factor of 2. The reason for this is that heavy elements have large positive charge on their nuclei, which strongly attract the electrons to the center of the atom. This contracts the size of the electron shells, so that more electrons fit in the only a slightly greater volume.
The temperature of a collection of atoms is a measure of the average energy of motion of those atoms; at 0 kelvin (absolute zero) atoms would have no motion. As the temperature of the system is increased, the kinetic energy of the particles in the system is increased, and their speed of motion increases. At room temperature, atoms making up gases in the air move at a speed of 500 m/s (about 1100 mph or 1800 km/h).
Elements, isotopes and ions
Atoms are generally classified by their atomic number Z, which corresponds to the number of protons in the atom. The atomic number determines which chemical element the atom is. For example, carbon atoms are atoms containing six protons. All atoms with the same atomic number share a wide variety of physical properties and exhibit the same chemical properties. The elements may be sorted according to the periodic table in order of increasing atomic number.
The atomic mass A, atomic mass number, or nucleon number of an element is the total number of protons and neutrons in an atom of that element, so-called because each proton and neutron has a mass of about 1 amu. The number of neutrons A−Z in an atom has no effect on which element it is. Each element can have numerous kinds of atoms with the same number of protons and electrons but varying numbers of neutrons. Each has the same atomic number but a different mass number. These are called the isotopes of an element. When writing the name of an isotope, the element name is followed by the mass number. For example, carbon-14 contains 6 protons and 8 neutrons in each atom, for a total mass number of 14.
The atomic mass listed for each element in the periodic table is an average of the isotope masses found in nature, weighted by their abundance.
The simplest atom is the hydrogen isotope protium, which has atomic number 1 and atomic mass number 1; it consists of one proton and one electron. The hydrogen isotope which also contains one neutron so is called deuterium or hydrogen-2; the hydrogen isotope with two neutrons is called tritium or hydrogen-3. Tritium is an unstable isotope which decays through a process called radioactivity. Almost all isotopes of each element are radioactive; only a few are stable. The elements with atomic number 84 (polonium) and heavier have no stable isotopes and are all radioactive.
Virtually all elements heavier than hydrogen and helium were created through stellar nucleosynthesis and supernova nucleosynthesis. Most of the elements lighter than uranium (Z=92) have stable-enough isotopes to occur naturally on Earth (with the notable exception of technetium Z=43). Several elements that do not occur on Earth have been found to be present in stars. Elements not normally found in nature have been artificially created by nuclear bombardment; as of 2006, elements have been created through atomic number 116 (given the temporary name ununhexium). These ultra-heavy elements are generally highly unstable and decay quickly.
Atoms that have either lost or gained electrons are called atomic ions (with either positive(+) or negative charge(−), respectively).
Valence and bonding
- see main article valence electrons and chemical bond
The number of electrons in an atom's outermost shell (the valence shell) governs its bonding behavior. Therefore, elements with the same number of valence electrons are grouped together in the columns of the periodic table of the elements. Alkali metals contain one electron on their outer shell; alkaline earth metals, two electrons; halogens, seven electrons; and various others.
Every atom is most stable with a full valence shell. This means that atoms with full valence shells (the noble gases) are very unreactive. Conversely, atoms with few electrons in their valence shell are more reactive it is. Alkali metals are therefore very reactive, with caesium, rubidium, and francium being the most reactive of all metals. Also, atoms that need only few electrons (such as the halogens) to fill their valence shells are reactive. Fluorine is the most reactive of all elements.
Atoms may fill their valence shells by chemical bonding. This can be achieved one of two ways: an atom can either share electrons with other atoms (a covalent bond), or it can remove electrons from (or donate electrons to) other atoms (an ionic bond). The formation of a bond causes a strong attraction between two atoms, creating molecules or ionic compounds. Many other types of bonds exist, including:
- polar covalent bonds;
- coordinate covalent bonds;
- metallic bonds;
- hydrogen bonds; and
- van der Waals bonds.
Atomic spectrum
- see main article Atomic spectroscopy
Since each element in the periodic table consists of an atom in a unique configuration with different numbers of protons and electrons, each element can also be uniquely described by the energies of its atomic orbitals and the number of electrons within them. Normally, an atom is found in its lowest-energy ground state; states with higher energy are called excited states. An electron may move from a lower-energy orbital to a higher-energy orbital by absorbing a photon with energy equal to the difference between the energies of the two levels. An electron in a higher-energy orbital may drop to a lower-energy orbital by emitting a photon. Since each element has a unique set of energy levels, each creates its own light pattern unique to itself: its own spectral signature.
If a set of atoms is heated (such as in an arc lamp), their electrons will move into excited states. When these atoms fall back toward the ground state, they will produce an emission spectrum. If a set of atoms is illuminated by a continuous spectrum, it will only absorb specific wavelengths (energies) of photon that correspond to the differences in its energy levels. The resulting pattern of gaps is called the absorption spectrum.
In spectroscopic analysis, scientists can use a spectrometer to study the atoms in stars and other distant objects. Due to the distinctive spectral lines that each element produces, they are able to tell the chemical composition of distant planets, stars and nebulae.
Not all parts of the atomic spectrum are in visible light part of the electromagnetic spectrum. For example, the hyperfine transitions (including the important 21 cm line) produce low-energy radio waves. When electrons deep inside large atoms are knocked out (for example by beta radiation), replacement atoms fall deep into the electric potential of the nucleus, producing high-energy x-rays.
Atoms and antimatter
- see main article antimatter
Antimatter can also form atoms, composed of positrons, antiprotons, and antineutrons. Since antimatter is very difficult to produce and store, only a small amount antihydrogen has ever existed on Earth. This was produced at CERN in the ATHENA and ATRAP experiments using the Antiproton Decelerator.
Atoms and the Big Bang
In models of the Big Bang, Big Bang nucleosynthesis predicts that within one to three minutes of the Big Bang almost all atomic material in the universe was created. During this process, nuclei of hydrogen and helium formed abundantly, but almost no elements heavier than lithium. Hydrogen makes up approximately 75% of the atoms in the universe; helium makes up 24%; and all other elements make up just 1%. However, although nuclei (fully-ionized atoms) were created, neutral atoms themselves could not form in the intense heat.
Big Bang chronology of the atom continues to approximately 379,000 years after the Big Bang when the cosmic temperature had dropped to just 3,000 K. It was then cool enough to allow the nuclei to capture electrons. This process is called recombination, during which the first neutral atoms took form. Once atoms become neutral, they only absorb photons of a discrete absorption spectrum. This allows most of the photons in the universe to travel unimpeded for billions of years. These photons are still detectable today in the cosmic microwave background.
After Big Bang nucleosynthesis, no heavier elements could be created until the formation of the first stars. These stars fused heavier elements through stellar nucleosynthesis during their lives and through supernova nucleosynthesis as they died. The seeding of the interstellar medium by heavy elements eventually allowed the formation of terrestrial planets like the Earth.
History of atomic theory
- Main article: Atomic theory
Philosophical atomism
From the 6th century BC, Hindu, Buddhist and Jaina philosophers in ancient India developed atomic theories (see Indian atomism). The first Indian philosopher who formulated ideas about the atom in a systematic manner was Kanada who lived in the 6th century BC. Another Indian philosopher, Pakudha Katyayana who also lived in the 6th century BC and was a contemporary of Gautama Buddha, had also propounded ideas about the atomic constitution of the material world.
Democritus and Leucippus, Greek philosophers in the 5th century BC, presented a theory of atoms. (See atomism for more details.) The Greeks believed that atoms were all made of the same material but had different shapes and sizes, which determined the physical properties of the material. For instance, the atoms of a liquid were thought to be smooth, allowing them to slide over each other.
None of these ideas, however, were founded in scientific experimentation.
Birth of modern atomic theory
In 1808, John Dalton proposed that an element is composed of atoms of a single, unique type, and that although their shape and structure was immutable, atoms of different elements could combine to form more complex structures (chemical compounds). He deduced this after the experimental discovery of the law of multiple proportions — that is, 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 whole numbers.
The experiment in question involved combining nitrous oxide (NO) with oxygen (O2). In one combination, these gases formed dinitrogen trioxide (N2O3), but when he repeated the combination with double the amount of oxygen (a ratio of 1:2), they instead formed nitrogen dioxide (NO2).
4NO + O2 → 2N2O3
4NO + 2O2 → 4NO2
Atomic theory conflicted with the theory of infinite divisibility, which states that matter can always be divided into smaller parts. In 1827, biologist Robert Brown observed that pollen grains floating in water constantly jiggled about for no apparent reason. In 1905, Albert Einstein theorised that this Brownian motion was caused by the water molecules continuously knocking the grains about, and developed a mathematical theory around it. This theory was validated experimentally in 1911 by French physicist Jean Perrin.
Discovery of subatomic particles
For much of this time, atoms were thought to be the smallest possible division of matter. However, in 1897, J.J. Thomson published his work proving that cathode rays are made of negatively charged particles (electrons). Since cathode rays are emitted from matter, this proved that atoms are made up of subatomic particles and are therefore divisible, and not the indivisible atomos postulated by Democritus. Physicists later invented a new term for such indivisible units, "elementary particles", since the word atom had come into its common modern use.
Study of atomic structure
At first, it was believed that the electrons were distributed more or less uniformly in a sea of positive charge (the plum pudding model). However, an experiment conducted in 1909 by colleagues of Ernest Rutherford demonstrated that atoms have a most of their mass and positive charge concentrated in a nucleus. In the gold foil experiment, alpha particles (emitted by polonium) were shot through a sheet of gold. Rutherford observed that most of the particles passed straight through the sheet with little deflection (striking a fluorescent screen on the other side). About 1 in 8000 of the alpha particles, however, were heavily deflected (by more than 90 degrees). This led to the planetary model of the atom in which pointlike electrons orbited in the space around a massive compact nucleus like planets orbiting the Sun.
The nucleus was later discovered to contain protons, and further experimentation by Rutherford found that the nuclear mass of most atoms surpassed that of the protons it possessed; this led him to postulate the existence of neutrons, whose existence would be proven in 1932 by James Chadwick.
The planetary model of the atom still had shortcomings. Firstly, a moving electric charge emits electromagnetic waves; according to classical electromagnetism, an orbiting charge would steadily lose energy and spiral towards the nucleus, colliding with it in a tiny fraction of a second. Secondly, the model did not explain why excited atoms emit light only in certain discrete spectra.
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 fixed amounts known as quanta. In 1913, Niels Bohr used this idea in 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. They were not allowed to spiral into the nucleus, because they could not lose energy in a continuous manner; they could only make quantum leaps between fixed energy levels. Bohr's model was extended by Arnold Sommerfeld in 1916 to include elliptical orbits, using a quantization of generalized momentum.
The ad hoc Bohr-Sommerfeld model was extremely difficult to use, but it made impressive predictions in agreement with certain spectral properties. However, the model was unable to explain multielectron atoms, predict transition rates or describe fine and hyperfine structure. In 1925, Erwin Schroedinger developed a full theory of quantum mechanics, described by the Schroedinger equation. Together with Wolfgang Pauli's exclusion principle, this allowed study of atoms with great precision when digital computers became available. Even today, these theories are used in the Hartree-Fock quantum chemical method to determine the energy levels of atoms. Further refinements of quantum theory such as the Dirac equation and quantum field theory made smaller impacts on the theory of atoms.
Another model of historical interest, proposed by Gilbert N. Lewis in 1916, had cubical atoms with electrons statically held at the corners. The cubes could share edges or faces to form chemical bonds. This model was created to account for chemical phenomena such as bonding, rather than physical phenomena such as atomic spectra.
See also
- Atomism
- Basic quantum mechanics
- Chemical bond
- Exotic atom
- Infinite divisibility
- List of particles
- Radioactive isotope
- Superatom
- Super-heavy atom
- Transuranium element
References
- Kenneth S. Krane, Introductory Nuclear Physics (1987)
External links
| Particles in physics - composite particles |
| Hadrons: Baryons (list) | Mesons (list)
Baryons: Nucleons | Hyperons | Exotic baryons | Pentaquarks |
