Metallicity

From Wikipedia, the free encyclopedia

The globular cluster M80. Stars in globular clusters are mainly older metal-poor members of Population II.

 

In astronomy, metallicity is used to describe the abundance of elements present in an object that are heavier than hydrogen or helium. Most of the physical matter in the Universe is in the form of hydrogen and helium, so astronomers use the word "metals" as a convenient short term for "all elements except hydrogen and helium". This usage is distinct from the usual physical definition of a solid metal. For example, stars and nebulae with relatively high abundances of carbon, nitrogen, oxygen, and neon are called "metal-rich" in astrophysical terms, even though those elements are non-metals in chemistry. 

The presence of heavier elements hails from stellar nucleosynthesis, the theory that the majority of elements heavier than hydrogen and helium in the Universe ("metals", hereafter) are formed in the cores of stars as they evolve. Over time, stellar winds and supernovae deposit the metals into the surrounding environment, enriching the interstellar medium and providing recycling materials for the birth of new stars. It follows that older generations of stars, which formed in the metal-poor early Universe, generally have lower metallicities than those of younger generations, which formed in a more metal-rich Universe. 

Observed changes in the chemical abundances of different types of stars, based on the spectral peculiarities that were later attributed to metallicity, led astronomer Walter Baade in 1944 to propose the existence of two different populations of stars. These became commonly known as Population I (metal-rich) and Population II (metal-poor) stars. A third stellar population was introduced in 1978, known as Population III stars. These extremely metal-poor stars were theorised to have been the "first-born" stars created in the Universe.

Common methods of calculation

Astronomers use several different methods to describe and approximate metal abundances, depending on the available tools and the object of interest. Some methods include determining the fraction of mass that is attributed to gas versus metals, or measuring the ratios of the number of atoms of two different elements as compared to the ratios found in the Sun.

Mass fraction

Stellar composition is often simply defined by the parameters X, Y and Z. Here X is the mass fraction of hydrogen, Y is the mass fraction of helium, and Z is the mass fraction of all the remaining chemical elements. Thus

${\displaystyle X+Y+Z=1.00.}$

In most stars, nebulae, H II regions, and other astronomical sources, hydrogen and helium are the two dominant elements. The hydrogen mass fraction is generally expressed as ${\displaystyle X\equiv m_{\text{H}}/M}$, where ${\displaystyle M}$ is the total mass of the system, and ${\displaystyle m_{\text{H}}}$ is the fractional mass of the hydrogen it contains. Similarly, the helium mass fraction is denoted as ${\displaystyle Y\equiv m_{\text{He}}/M}$. The remainder of the elements are collectively referred to as "metals", and the metallicity—the mass fraction of elements heavier than helium—can be calculated as

${\displaystyle Z=\sum _{i>{\text{He}}}{\frac {m_{i}}{M}}=1-X-Y.}$

For the surface of the Sun, these parameters are measured to have the following values:

Description	Solar value
Hydrogen mass fraction	${\displaystyle X_{\text{sun}}=0.7381}$
Helium mass fraction	${\displaystyle Y_{\text{sun}}=0.2485}$
Metallicity	${\displaystyle Z_{\text{sun}}=0.0134}$

Due to the effects of stellar evolution, neither the initial composition nor the present day bulk composition of the Sun is the same as its present-day surface composition.

Chemical abundance ratios

The overall stellar metallicity is often defined using the total iron content of the star, as iron is among the easiest to measure with spectral observations in the visible spectrum (even though oxygen is the most abundant heavy element – see metallicities in HII regions below). The abundance ratio is defined as the logarithm of the ratio of a star's iron abundance compared to that of the Sun and is expressed thus:

${\displaystyle [{\text{Fe}}/{\text{H}}]=\log _{10}{\left({\frac {N_{\text{Fe}}}{N_{\text{H}}}}\right)_{\text{star}}}-\log _{10}{\left({\frac {N_{\text{Fe}}}{N_{\text{H}}}}\right)_{\text{sun}}},}$

where ${\displaystyle N_{\text{Fe}}}$ and ${\displaystyle N_{\text{H}}}$ are the number of iron and hydrogen atoms per unit of volume respectively. The unit often used for metallicity is the dex, contraction of "decimal exponent". By this formulation, stars with a higher metallicity than the Sun have a positive logarithmic value, whereas those with a lower metallicity than the Sun have a negative value. For example, stars with a [Fe/H] value of +1 have 10 times the metallicity of the Sun (10¹); conversely, those with a [Fe/H] value of −1 have 1/10, while those with a [Fe/H] value of 0 have the same metallicity as the Sun, and so on. Young Population I stars have significantly higher iron-to-hydrogen ratios than older Population II stars. Primordial Population III stars are estimated to have a metallicity of less than −6.0, that is, less than a millionth of the abundance of iron in the Sun.

The same notation is used to express variations in abundances between other the individual elements as compared to solar proportions. For example, the notation "[O/Fe]" represents the difference in the logarithm of the star's oxygen abundance versus its iron content compared to that of the Sun. In general, a given stellar nucleosynthetic process alters the proportions of only a few elements or isotopes, so a star or gas sample with nonzero [X/Fe] values may be showing the signature of particular nuclear processes.

Photometric colors

Astronomers can estimate metallicities through measured and calibrated systems that correlate photometric measurements and spectroscopic measurements. For example, the Johnson UVB filters can be used to detect an ultraviolet (UV) excess in stars, where a larger UV excess indicates a larger presence of metals that absorb the UV radiation, thereby making the star appear "redder". The UV excess, δ(U−B), is defined as the difference between a star's U and B band magnitudes, compared to the difference between U and B band magnitudes of metal-rich stars in the Hyades cluster. Unfortunately, δ(U−B) is sensitive to both metallicity and temperature: if two stars are equally metal-rich, but one is cooler than the other, they will likely have different δ(U−B) values. To help mitigate this degeneracy, a star's B−V color can be used as an indicator for temperature. Furthermore, the UV excess and B−V color can be corrected to relate the δ(U−B) value to iron abundances.

Other photometric systems that can be used to determine metallicities of certain astrophysical objects include the Strӧmgren system, the Geneva system, the Washington system, and the DDO system.

Metallicities in various astrophysical objects

Stars

Relationship between stellar metallicity and planets

A star's metallicity measurement is one parameter that helps determine whether a star has planets and the type of planets, as there is a direct correlation between metallicity and the type of planets a star may have. Measurements have demonstrated the connection between a star's metallicity and gas giant planets, like Jupiter and Saturn. The more metals in a star and thus its planetary system and proplyd, the more likely the system may have gas giant planets and rocky planets. Current models show that the metallicity along with the correct planetary system temperature and distance from the star are key to planet and planetesimal formation. For two stars that have equal age and mass but different metallicity, the less metallic star is bluer. Among stars of the same color, less metallic stars emit more ultraviolet radiation. The Sun, with 8 planets and 5 known dwarf planets, is used as the reference, with a [Fe/H] of 0.00.

HII regions

Young, massive and hot stars (typically of spectral types O and B) in H II regions emit UV photons that ionize ground-state hydrogen atoms, knocking electrons and protons free; this process is known as photoionization. The free electrons can strike other atoms nearby, exciting bound metallic electrons into a metastable state, which eventually decay back into a ground state, emitting photons with energies that correspond to forbidden lines. Through these transitions, astronomers have developed several observational methods to estimate metal abundances in HII regions, where the stronger the forbidden lines in spectroscopic observations, the higher the metallicity. These methods are dependent on one or more of the following: the variety of asymmetrical densities inside HII regions, the varied temperatures of the embedded stars, and/or the electron density within the ionized region.

Theoretically, to determine the total abundance of a single element in an HII region, all transition lines should be observed and summed. However, this can be observationally difficult due to variation in line strength. Some of the most common forbidden lines used to determine metal abundances in HII regions are from oxygen (e.g. [O II] λ = (3727, 7318, 7324) Å, and [O III] λ = (4363, 4959, 5007) Å, nitrogen (e.g. [NII] λ = (5755, 6548, 6584) Å), and sulfur (e.g. [SII] λ = (6717,6731) Å and [SIII] λ = (6312, 9069, 9531) Å) in the optical spectrum, and the [OIII] λ = (52, 88) μm and [NIII] λ = 57 μm lines in the infrared spectrum. Oxygen has some of the stronger, more abundant lines in HII regions, making it a main target for metallicity estimates within these objects. To calculate metal abundances in HII regions using oxygen flux measurements, astronomers often use the R₂₃ method, in which

${\displaystyle R_{23}={\frac {[{\text{O II}}]_{3727~\mathrm {\AA} }+[{\text{O III}}]_{4959~\mathrm {\AA} +5007~\mathrm {\AA} }}{{\text{H}}_{\beta }}},}$

where ${\displaystyle {\text{O III}}_{3727~\mathrm {\AA} }+{\text{O III}}_{4959~\mathrm {\AA} +5007~\mathrm {\AA} }}$ is the sum of the fluxes from oxygen emission lines measured at the rest frame λ = (3727, 4959 and 5007) Å wavelengths, divided by the flux from the H_β emission line at the rest frame λ = 4861 Å wavelength. This ratio is well defined through models and observational studies, but caution should be taken, as the ratio is often degenerate, providing both a low and high metallicity solution, which can be broken with additional line measurements. Similarly, other strong forbidden line ratios can be used, e.g. for sulfur, where

${\displaystyle S_{23}={\frac {[{\text{S II}}]_{6716~\mathrm {\AA} +6731~\mathrm {\AA} }+[{\text{S III}}]_{9069~\mathrm {\AA} +9532~\mathrm {\AA} }}{{\text{H}}_{\beta }}}.}$

Metal abundances within HII regions are typically less than 1%, with the percentage decreasing on average with distance from the Galactic Center.

Big Bang nucleosynthesis

From Wikipedia, the free encyclopedia

In physical cosmology, Big Bang nucleosynthesis (abbreviated BBN, also known as primordial nucleosynthesis, arch(a)eonucleosynthesis, archonucleosynthesis, protonucleosynthesis and pal(a)eonucleosynthesis) refers to the production of nuclei other than those of the lightest isotope of hydrogen (hydrogen-1, ¹H, having a single proton as a nucleus) during the early phases of the Universe. Primordial nucleosynthesis is believed by most cosmologists to have taken place in the interval from roughly 10 seconds to 20 minutes after the Big Bang, and is calculated to be responsible for the formation of most of the universe's helium as the isotope helium-4 (⁴He), along with small amounts of the hydrogen isotope deuterium (²H or D), the helium isotope helium-3 (³He), and a very small amount of the lithium isotope lithium-7 (⁷Li). In addition to these stable nuclei, two unstable or radioactive isotopes were also produced: the heavy hydrogen isotope tritium (³H or T); and the beryllium isotope beryllium-7 (⁷Be); but these unstable isotopes later decayed into ³He and ⁷Li, as above. 

Essentially all of the elements that are heavier than lithium were created much later, by stellar nucleosynthesis in evolving and exploding stars.

Characteristics

There are several important characteristics of Big Bang nucleosynthesis (BBN):

• The initial conditions (neutron-proton ratio) were set in the first second after the Big Bang.
• The universe was very close to homogeneous at this time, and strongly radiation-dominated.
• The fusion of nuclei occurred between roughly 10 seconds to 20 minutes after the Big Bang; this corresponds to the temperature range when the universe was cool enough for deuterium to survive, but hot and dense enough for fusion reactions to occur at a significant rate.
• It was widespread, encompassing the entire observable universe.

The key parameter which allows one to calculate the effects of BBN is the baryon/photon number ratio, which is a small number of order 6 × 10⁻¹⁰. This parameter corresponds to the baryon density and controls the rate at which nucleons collide and react; from this it is possible to calculate element abundances after nucleosynthesis ends. Although the baryon per photon ratio is important in determining element abundances, the precise value makes little difference to the overall picture. Without major changes to the Big Bang theory itself, BBN will result in mass abundances of about 75% of hydrogen-1, about 25% helium-4, about 0.01% of deuterium and helium-3, trace amounts (on the order of 10⁻¹⁰) of lithium, and negligible heavier elements. That the observed abundances in the universe are generally consistent with these abundance numbers is considered strong evidence for the Big Bang theory. 

In this field, for historical reasons it is customary to quote the helium-4 fraction by mass, symbol Y, so that 25% helium-4 means that helium-4 atoms account for 25% of the mass, but less than 8% of the nuclei would be helium-4 nuclei. Other (trace) nuclei are usually expressed as number ratios to hydrogen.

Important parameters

The creation of light elements during BBN was dependent on a number of parameters; among those was the neutron-proton ratio (calculable from Standard Model physics) and the baryon-photon ratio.

Neutron–proton ratio

The neutron-proton ratio was set by Standard Model physics before the nucleosynthesis era, essentially within the first 1-second after the Big Bang. Neutrons can react with positrons or electron neutrinos to create protons and other products in one of the following reactions:

${\displaystyle {\ce {n\ +e+<=>{\overline {\nu }}_{e}+p}}}$

${\displaystyle {\ce {n \ + \nu_{e}<=> p + e-}}}$

At times much earlier than 1 sec, these reactions were fast and maintained the n/p ratio close to 1:1. As the temperature dropped, the equilibrium shifted in favour of protons due to their slightly lower mass, and the n/p ratio smoothly decreased. These reactions continued until the decreasing temperature and density caused the reactions to become too slow, which occurred at about T = 0.7 MeV (time around 1 second) and is called the freeze out temperature. At freeze out, the neutron-proton ratio was about 1/6. However, free neutrons are unstable with a mean life of 880 sec; some neutrons decayed in the next few minutes before fusing into any nucleus, so the ratio of total neutrons to protons after nucleosynthesis ends is about 1/7. Almost all neutrons that fused instead of decaying ended up combined into helium-4, due to the fact that helium-4 has the highest binding energy per nucleon among light elements. This predicts that about 8% of all atoms should be helium-4, leading to a mass fraction of helium-4 of about 25%, which is in line with observations. Small traces of deuterium and helium-3 remained as there was insufficient time and density for them to react and form helium-4.

Baryon–photon ratio

The baryon–photon ratio, η, is the key parameter determining the abundances of light elements after nucleosynthesis ends. Baryons and light elements can fuse in the following main reactions:

${\displaystyle {\ce {p + n -> ^2H + \gamma}}}$ ${\displaystyle {\ce {p + ^2H -> ^3He + \gamma}}}$ ${\displaystyle {\ce {^2H + ^2H -> ^3He + n}}}$ ${\displaystyle {\ce {^2H + ^2H -> ^3H + p}}}$ ${\displaystyle {\ce {^3He + ^2H -> ^4He + p}}}$ ${\displaystyle {\ce {^3H + ^2H -> ^4He + n}}}$

along with some other low-probability reactions leading to ⁷Li or ⁷Be. (An important feature is that there are no stable nuclei with mass 5 or 8, which implies that reactions adding one baryon to ⁴He, or fusing two ⁴He, do not occur). Most fusion chains during BBN ultimately terminate in ⁴He (helium-4), while "incomplete" reaction chains lead to small amounts of left-over ²H or ³He; the amount of these decreases with increasing baryon-photon ratio. That is, the larger the baryon-photon ratio the more reactions there will be and the more efficiently deuterium will be eventually transformed into helium-4. This result makes deuterium a very useful tool in measuring the baryon-to-photon ratio.

Sequence

The main nuclear reaction chains for Big Bang nucleosynthesis

 

Big Bang nucleosynthesis began roughly 10 seconds after the big bang, when the universe had cooled sufficiently to allow deuterium nuclei to survive disruption by high-energy photons. (Note that the neutron-proton freeze-out time was earlier). This time is essentially independent of dark matter content, since the universe was highly radiation dominated until much later, and this dominant component controls the temperature/time relation. At this time there were about six protons for every neutron, but a small fraction of the neutrons decay before fusing in the next few hundred seconds, so at the end of nucleosynthesis there are about seven protons to every neutron, and almost all the neutrons are in Helium-4 nuclei. The sequence of these reaction chains is shown on the image.

One feature of BBN is that the physical laws and constants that govern the behavior of matter at these energies are very well understood, and hence BBN lacks some of the speculative uncertainties that characterize earlier periods in the life of the universe. Another feature is that the process of nucleosynthesis is determined by conditions at the start of this phase of the life of the universe, and proceeds independently of what happened before. 

As the universe expands, it cools. Free neutrons are less stable than helium nuclei, and the protons and neutrons have a strong tendency to form helium-4. However, forming helium-4 requires the intermediate step of forming deuterium. Before nucleosynthesis began, the temperature was high enough for many photons to have energy greater than the binding energy of deuterium; therefore any deuterium that was formed was immediately destroyed (a situation known as the deuterium bottleneck). Hence, the formation of helium-4 is delayed until the universe became cool enough for deuterium to survive (at about T = 0.1 MeV); after which there was a sudden burst of element formation. However, very shortly thereafter, around twenty minutes after the Big Bang, the temperature and density became too low for any significant fusion to occur. At this point, the elemental abundances were nearly fixed, and the only changes were the result of the radioactive decay of the two major unstable products of BBN, tritium and beryllium-7.

History of theory

The history of Big Bang nucleosynthesis began with the calculations of Ralph Alpher in the 1940s. Alpher published the Alpher–Bethe–Gamow paper that outlined the theory of light-element production in the early universe. 

During the 1970s, there was a major puzzle in that the density of baryons as calculated by Big Bang nucleosynthesis was much less than the observed mass of the universe based on measurements of galaxy rotation curves and galaxy cluster dynamics. This puzzle was resolved in large part by postulating the existence of dark matter.

Heavy elements

A version of the periodic table indicating the origins – including big bang nucleosynthesis – of the elements. All elements above 103 (lawrencium) are also manmade and are not included.

 

Big Bang nucleosynthesis produced very few nuclei of elements heavier than lithium due to a bottleneck: the absence of a stable nucleus with 8 or 5 nucleons. This deficit of larger atoms also limited the amounts of lithium-7 produced during BBN. In stars, the bottleneck is passed by triple collisions of helium-4 nuclei, producing carbon (the triple-alpha process). However, this process is very slow and requires much higher densities, taking tens of thousands of years to convert a significant amount of helium to carbon in stars, and therefore it made a negligible contribution in the minutes following the Big Bang.

The predicted abundance of CNO isotopes produced in Big Bang nucleosynthesis is expected to be on the order of 10⁻¹⁵ that of H, making them essentially undetectable and negligible. Indeed, none of these primordial isotopes of the elements from lithium to oxygen have yet been detected, although those of beryllium and boron may be able to be detected in the future. So far, the only stable nuclides known experimentally to have been made before or during Big Bang nucleosynthesis are protium, deuterium, helium-3, helium-4, and lithium-7.

Helium-4

Big Bang nucleosynthesis predicts a primordial abundance of about 25% helium-4 by mass, irrespective of the initial conditions of the universe. As long as the universe was hot enough for protons and neutrons to transform into each other easily, their ratio, determined solely by their relative masses, was about 1 neutron to 7 protons (allowing for some decay of neutrons into protons). Once it was cool enough, the neutrons quickly bound with an equal number of protons to form first deuterium, then helium-4. Helium-4 is very stable and is nearly the end of this chain if it runs for only a short time, since helium neither decays nor combines easily to form heavier nuclei (since there are no stable nuclei with mass numbers of 5 or 8, helium does not combine easily with either protons, or with itself). Once temperatures are lowered, out of every 16 nucleons (2 neutrons and 14 protons), 4 of these (25% of the total particles and total mass) combine quickly into one helium-4 nucleus. This produces one helium for every 12 hydrogens, resulting in a universe that is a little over 8% helium by number of atoms, and 25% helium by mass. 

One analogy is to think of helium-4 as ash, and the amount of ash that one forms when one completely burns a piece of wood is insensitive to how one burns it. The resort to the BBN theory of the helium-4 abundance is necessary as there is far more helium-4 in the universe than can be explained by stellar nucleosynthesis. In addition, it provides an important test for the Big Bang theory. If the observed helium abundance is significantly different from 25%, then this would pose a serious challenge to the theory. This would particularly be the case if the early helium-4 abundance was much smaller than 25% because it is hard to destroy helium-4. For a few years during the mid-1990s, observations suggested that this might be the case, causing astrophysicists to talk about a Big Bang nucleosynthetic crisis, but further observations were consistent with the Big Bang theory.

Deuterium

Deuterium is in some ways the opposite of helium-4, in that while helium-4 is very stable and difficult to destroy, deuterium is only marginally stable and easy to destroy. The temperatures, time, and densities were sufficient to combine a substantial fraction of the deuterium nuclei to form helium-4 but insufficient to carry the process further using helium-4 in the next fusion step. BBN did not convert all of the deuterium in the universe to helium-4 due to the expansion that cooled the universe and reduced the density, and so cut that conversion short before it could proceed any further. One consequence of this is that, unlike helium-4, the amount of deuterium is very sensitive to initial conditions. The denser the initial universe was, the more deuterium would be converted to helium-4 before time ran out, and the less deuterium would remain. 

There are no known post-Big Bang processes which can produce significant amounts of deuterium. Hence observations about deuterium abundance suggest that the universe is not infinitely old, which is in accordance with the Big Bang theory. 

During the 1970s, there were major efforts to find processes that could produce deuterium, but those revealed ways of producing isotopes other than deuterium. The problem was that while the concentration of deuterium in the universe is consistent with the Big Bang model as a whole, it is too high to be consistent with a model that presumes that most of the universe is composed of protons and neutrons. If one assumes that all of the universe consists of protons and neutrons, the density of the universe is such that much of the currently observed deuterium would have been burned into helium-4. The standard explanation now used for the abundance of deuterium is that the universe does not consist mostly of baryons, but that non-baryonic matter (also known as dark matter) makes up most of the mass of the universe. This explanation is also consistent with calculations that show that a universe made mostly of protons and neutrons would be far more clumpy than is observed.

It is very hard to come up with another process that would produce deuterium other than by nuclear fusion. Such a process would require that the temperature be hot enough to produce deuterium, but not hot enough to produce helium-4, and that this process should immediately cool to non-nuclear temperatures after no more than a few minutes. It would also be necessary for the deuterium to be swept away before it reoccurs.

Producing deuterium by fission is also difficult. The problem here again is that deuterium is very unlikely due to nuclear processes, and that collisions between atomic nuclei are likely to result either in the fusion of the nuclei, or in the release of free neutrons or alpha particles. During the 1970s, cosmic ray spallation was proposed as a source of deuterium. That theory failed to account for the abundance of deuterium, but led to explanations of the source of other light elements.

Lithium

Lithium-7 and lithium-6 produced in the Big Bang are in the order of: lithium-7 to be 10⁻⁹ of all primordial nuclides; and lithium-6 around 10⁻¹³.

Measurements and status of theory

The theory of BBN gives a detailed mathematical description of the production of the light "elements" deuterium, helium-3, helium-4, and lithium-7. Specifically, the theory yields precise quantitative predictions for the mixture of these elements, that is, the primordial abundances at the end of the big-bang. 

In order to test these predictions, it is necessary to reconstruct the primordial abundances as faithfully as possible, for instance by observing astronomical objects in which very little stellar nucleosynthesis has taken place (such as certain dwarf galaxies) or by observing objects that are very far away, and thus can be seen in a very early stage of their evolution (such as distant quasars). 

As noted above, in the standard picture of BBN, all of the light element abundances depend on the amount of ordinary matter (baryons) relative to radiation (photons). Since the universe is presumed to be homogeneous, it has one unique value of the baryon-to-photon ratio. For a long time, this meant that to test BBN theory against observations one had to ask: can all of the light element observations be explained with a single value of the baryon-to-photon ratio? Or more precisely, allowing for the finite precision of both the predictions and the observations, one asks: is there some range of baryon-to-photon values which can account for all of the observations?

More recently, the question has changed: Precision observations of the cosmic microwave background radiation with the Wilkinson Microwave Anisotropy Probe (WMAP) and Planck give an independent value for the baryon-to-photon ratio. Using this value, are the BBN predictions for the abundances of light elements in agreement with the observations?

The present measurement of helium-4 indicates good agreement, and yet better agreement for helium-3. But for lithium-7, there is a significant discrepancy between BBN and WMAP/Planck, and the abundance derived from Population II stars. The discrepancy is a factor of 2.4―4.3 below the theoretically predicted value and is considered a problem for the original models,^[14] that have resulted in revised calculations of the standard BBN based on new nuclear data, and to various reevaluation proposals for primordial proton-proton nuclear reactions, especially the abundances of ⁷Be + n → ⁷Li + p, versus ⁷Be + ²H → ⁸Be + p.

Non-standard scenarios

In addition to the standard BBN scenario there are numerous non-standard BBN scenarios. These should not be confused with non-standard cosmology: a non-standard BBN scenario assumes that the Big Bang occurred, but inserts additional physics in order to see how this affects elemental abundances. These pieces of additional physics include relaxing or removing the assumption of homogeneity, or inserting new particles such as massive neutrinos.

There have been, and continue to be, various reasons for researching non-standard BBN. The first, which is largely of historical interest, is to resolve inconsistencies between BBN predictions and observations. This has proved to be of limited usefulness in that the inconsistencies were resolved by better observations, and in most cases trying to change BBN resulted in abundances that were more inconsistent with observations rather than less. The second reason for researching non-standard BBN, and largely the focus of non-standard BBN in the early 21st century, is to use BBN to place limits on unknown or speculative physics. For example, standard BBN assumes that no exotic hypothetical particles were involved in BBN. One can insert a hypothetical particle (such as a massive neutrino) and see what has to happen before BBN predicts abundances that are very different from observations. This has been done to put limits on the mass of a stable tau neutrino.

History of the periodic table

From Wikipedia, the free encyclopedia

Dalton (1806): listing the known elements by atomic weight

 

Mendeleev (1871): tabular ordering, showing periodic behavior

 

2016: current form, 118 known elements

 

The periodic table is an arrangement of the chemical elements, which are organized on the basis of their atomic numbers, electron configurations and recurring chemical properties. Elements are presented in order of increasing atomic number. The standard form of the table consists of a grid with rows called periods and columns called groups. 

The history of the periodic table reflects over two centuries of growth in the understanding of chemical properties, with major contributions made by Antoine-Laurent de Lavoisier, Johann Wolfgang Döbereiner, John Newlands, Julius Lothar Meyer, Dmitri Mendeleev, and Glenn T. Seaborg.

Antiquity to the 18th century

A number of physical elements (such as platinum, mercury, tin and zinc) have been known from antiquity, as they are found in their native form and are relatively simple to mine with primitive tools. Around 330 BCE, the Greek philosopher Aristotle proposed that everything is made up of a mixture of one or more roots, an idea that had originally been suggested by the Sicilian philosopher Empedocles. The four roots, which were later renamed as elements by Plato, were earth, water, air and fire. Similar ideas about these four elements also existed in other ancient traditions, such as Indian philosophy.

Hennig Brand, as shown in The Alchemist Discovering Phosphorus

Hennig Brand

The history of the periodic table is also a history of the discovery of the chemical elements. The first person in history to discover a new element was Hennig Brand, a bankrupt German merchant. Brand tried to discover the Philosopher's Stone — a mythical object that was supposed to turn inexpensive base metals into gold. In 1669 (or later), his experiments with distilled human urine resulted in the production of a glowing white substance, which he called "cold fire" (kaltes Feuer). He kept his discovery secret until 1680, when Robert Boyle rediscovered phosphorus and published his findings. The discovery of phosphorus helped to raise the question of what it meant for a substance to be an element. 

In 1661, Boyle defined an element as "those primitive and simple Bodies of which the mixt ones are said to be composed, and into which they are ultimately resolved."

Antoine-Laurent de Lavoisier

Antoine Laurent de Lavoisier

 

Lavoisier's Traité Élémentaire de Chimie (Elementary Treatise of Chemistry), which was written in 1789 and first translated into English by the writer Robert Kerr, is considered to be the first modern textbook about chemistry. Lavoisier defined an element as a substance that cannot be broken down into a simpler substance by a chemical reaction. This simple definition served for a century and lasted until the discovery of subatomic particles. Lavoisier's book contained a list of "simple substances" that Lavoisier believed could not be broken down further, which included oxygen, nitrogen, hydrogen, phosphorus, mercury, zinc and sulfur, which formed the basis for the modern list of elements. Lavoisier's list also included 'light' and 'caloric', which at the time were believed to be material substances. He classified these substances into metals and non metals. While many leading chemists refused to believe Lavoisier's new revelations, the Elementary Treatise was written well enough to convince the younger generation. However, Lavoisier's descriptions of his elements lack completeness, as he only classified them as metals and non-metals.

19th century

William Prout

In 1815, the English physician and chemist William Prout noticed that atomic weights seemed to be multiples of that of hydrogen.

Johann Wolfgang Döbereiner

In 1817, Johann Wolfgang Döbereiner, a chemist, began to formulate one of the earliest attempts to classify the elements. In 1829, he found that he could form some of the elements into groups of three, with the members of each group having related properties. He termed these groups triads.

Definition of Triad law:-"Chemically analogous elements arranged in increasing order of their atomic weights formed well marked groups of three called Triads in which the atomic weight of the middle element was found to be generally the arithmetic mean of the atomic weight of the other two elements in the triad.

1. chlorine, bromine, and iodine
2. calcium, strontium, and barium
3. sulfur, selenium, and tellurium
4. lithium, sodium, and potassium

Alexandre-Emile Béguyer de Chancourtois

Alexandre-Emile Béguyer de Chancourtois, a French geologist, was the first person to notice the periodicity of the elements — similar elements occurring at regular intervals when they are ordered by their atomic weights. In 1862 he devised an early form of periodic table, which he named Vis tellurique (the 'telluric helix'), after the element tellurium, which fell near the center of his diagram. With the elements arranged in a spiral on a cylinder by order of increasing atomic weight, de Chancourtois saw that elements with similar properties lined up vertically. His 1863 publication included a chart (which contained ions and compounds, in addition to elements), but his original paper in the Comptes Rendus de l'Académie des Sciences used geological rather than chemical terms and did not include a diagram. As a result, de Chancourtois' ideas received little attention until after the work of Dmitri Mendeleev had been published.

John Newlands

Newlands' law of octaves

 

In 1864, the English chemist John Newlands classified the sixty-two known elements into eight groups, based on their physical properties.

Newlands noted that many pairs of similar elements existed, which differed by some multiple of eight in mass number, and was the first to assign them an atomic number. When his 'law of octaves' was printed in Chemistry News, likening this periodicity of eights to the musical scale, it was ridiculed by some of his contemporaries. His lecture to the Chemistry Society on 1 March 1866 was not published, the Society defending their decision by saying that such 'theoretical' topics might be controversial.

The importance of Newlands' analysis was eventually recognised by the Chemistry Society with a Gold Medal five years after they recognised Mendeleev's work. It was not until the following century, with Gilbert N. Lewis's valence bond theory (1916) and Irving Langmuir's octet theory of chemical bonding (1919), that the importance of the periodicity of eight would be accepted. The Royal Chemistry Society acknowledged Newlands' contribution to science in 2008, when they put a Blue Plaque on the house where he was born, which described him as the "discoverer of the Periodic Law for the chemical elements".

He contributed the word 'periodic' in chemistry.

Julius Lothar Meyer

Julius Lothar Meyer's periodic table, published in "Die modernen Theorien der Chemie" (1864)

 

Meyer noted, as J. A. R. Newlands did in England, if each element is arranged in the order of their atomic weights, they fall into groups of similar chemical and physical properties repeated at periodic intervals. According to him, if the atomic weights were plotted as ordinates and the atomic volumes as abscissae—the curve obtained a series of maxima and minima—the most electro-positive elements appearing at the peaks of the curve in the order of their atomic weights. 

His book, Die modernen Theorien der Chemie, which he began writing in Breslau in 1862 and which was published two years later, contained an early version of the periodic table containing 28 elements, classified elements into six families by their valence—for the first time, elements had been grouped according to their valence. Works on organizing the elements by atomic weight, until then had been stymied by inaccurate measurements of the atomic weights.

He published articles about classification table of the elements in horizontal form (1862, 1864) and vertical form (1870), in which the series of periods are properly ended by an element of the alkaline earth metal group.

In 1869, a few months later than Mendeleev, Meyer published a revised and expanded version of his 1864 table independently, which was similar to that published by Mendeleev (Meyer had been sent a copy of Mendeleev's table earlier; Mendeleev had sent it to many well-known chemists of his day) and a paper showing graphically the periodicity of the elements as a function of atomic weight. 

In 1882, both Meyer and Mendeleev received the Davy Medal from the Royal Society in recognition of their work on the Periodic Law.

Dmitri Mendeleev

Dmitri Ivanovich Mendeleev

 

Zeitschrift für Chemie (1869, pages 405–6), in which Mendeleev's periodic table is first published outside Russia.

 

Mendeleev's 1871 periodic table. Dashes: unknown elements. Group I-VII: modern group 1–2 and 3–7 with transition metals added; some of these extend into a group VIII. Noble gases unknown (and unpredicted).

 

The Russian chemist Dmitri Mendeleev arranged the elements by atomic mass, corresponding to relative molar mass. It is sometimes said that he played 'chemical solitaire' on long train journeys, using cards with various facts about the known elements. On March 1 [O.S. February 17] 1869, he put a date on his first table, and sent it for publication. On 18 March [O.S. 6 March] 1869, Mendeleev gave a formal presentation, The Dependence Between the Properties of the Atomic Weights of the Elements, to the Russian Chemical Society. In 1869, the table was published in an obscure Russian journal and then republished in a German journal, Zeitschrift für Chemie. In it, Mendeleev stated that:

1. The elements, if arranged according to their atomic mass, exhibit an apparent periodicity of properties.
2. Elements which are similar as regards to their chemical properties have atomic weights which are either of nearly the same value (e.g., Pt, Ir, Os) or which increase regularly (e.g., K, Rb, Cs).
3. The arrangement of the elements, or of groups of elements in the order of their atomic masses, corresponds to their so-called valencies, as well as, to some extent, to their distinctive chemical properties; as is apparent among other series in that of Li, Be, B, C, N, O, and F.
4. The elements which are the most widely diffused have small atomic weights.
5. The magnitude of the atomic weight determines the character of the element, just as the magnitude of the molecule determines the character of a compound body.
6. We must expect the discovery of many yet unknown elements – for example, elements analogous to aluminium and silicon – whose atomic weight would be between 65 and 75.
7. The atomic weight of an element may sometimes be amended by a knowledge of those of its contiguous elements. Thus the atomic weight of tellurium must lie between 123 and 126, and cannot be 128.
8. Certain characteristic properties of elements can be foretold from their atomic masses.

Scientific benefits of Mendeleev's table

• It enabled Mendeleev to predict the discovery of new elements and left spaces for them, namely eka-silicon (germanium, discovered in 1885), eka-aluminium (gallium, 1875), and eka-boron (scandium, 1879). Thus, there was no disturbance in the periodic table.
• It could be used by Mendeleev to point out that some of the atomic weights being used at the time were incorrect.
• It provided for variance from atomic weight order.

William Odling

In 1864, the English chemist William Odling also drew up a table that was remarkably similar to the table produced by Mendeleev. Odling overcame the tellurium-iodine problem and even managed to get thallium, lead, mercury and platinum into the right groups, which is something that Mendeleev failed to do at his first attempt. Odling failed to achieve recognition, however, since it is suspected that he, as Secretary of the Chemical Society of London, was instrumental in discrediting Newlands' earlier work on the periodic table.

Shortcomings of early versions of the periodic table

• The table was not able to predict the existence of the noble gases, but did leave spaces for yet-to-be discovered elements. Time proved this method correct. When the entire group of noble gases was discovered, primarily by William Ramsay, he added them to the table as Group 0, without disturbing the basic concept of the periodic table.
• A single position could not be assigned to hydrogen, which could be placed either in the alkali metals group, the halogens group or separately above the table between boron and carbon.
• The lanthanides were difficult to fit into the table.

20th century

Frederick Soddy

By 1912 almost 50 different radioactive elements had been found, too many for the periodic table. Frederick Soddy in 1913 found that although they emitted different radiation, many elements were alike in their chemical characteristics so shared the same place on the table. They became known as isotopes, from the Greek eisos topos ("same place").

Henry Moseley

Henry Moseley

 

In 1914, a year before he was killed in action at Gallipoli, the English physicist Henry Moseley found a relationship between the X-ray wavelength of an element and its atomic number. He was then able to re-sequence the periodic table by nuclear charge, rather than by atomic weight. Before this discovery, atomic numbers were sequential numbers based on an element's atomic weight. Moseley's discovery showed that atomic numbers were in fact based upon experimental measurements. 

Using information about their X-ray wavelengths, Moseley placed argon (with an atomic number Z=18) before potassium (Z=19), despite the fact that argon's atomic weight of 39.9 is greater than the atomic weight of potassium (39.1). The new order was in agreement with the chemical properties of these elements, since argon is a noble gas and potassium is an alkali metal. Similarly, Moseley placed cobalt before nickel and was able to explain that tellurium occurs before iodine, without revising the experimental atomic weight of tellurium, as had been proposed by Mendeleev.

Moseley's research showed that there were gaps in the periodic table at atomic numbers 43 and 61, which are now known to be occupied by technetium and promethium respectively.

Glenn T. Seaborg

During his Manhattan Project research in 1943, Glenn T. Seaborg experienced unexpected difficulties in isolating the elements americium and curium. These elements, in addition to the elements from actinium to plutonium, were believed to form a fourth series of transition metals. Seaborg wondered if these elements belonged to a different series, which would explain why their chemical properties, in particular the instability of higher oxidation states, were different from predictions. In 1945, against the advice of colleagues, he proposed a significant change to Mendeleev's table: the actinide series.

Seaborg's actinide concept of heavy element electronic structure, predicting that the actinides form a transition series analogous to the rare earth series of lanthanide elements, is now well accepted and included in the periodic table. The actinide series is the second row of the f-block (5f series). In both the actinide and lanthanide series, an inner electron shell is being filled. The actinide series comprises the elements from actinium to lawrencium. Seaborg's subsequent elaborations of the actinide concept theorized a series of superheavy elements in a transactinide series comprising elements from 104 to 121 and a superactinide series of elements from 122 to 153.