Gemstone

From Wikipedia, the free encyclopedia

https://en.wikipedia.org/wiki/Gemstone

Group of precious and semiprecious stones—both uncut and faceted—including (clockwise from top left) diamond, uncut synthetic sapphire, ruby, uncut emerald, and amethyst crystal cluster.

A gemstone (also called a fine gem, jewel, precious stone, semiprecious stone, or simply gem) is a piece of mineral crystal which, when cut or polished, is used to make jewelry or other adornments. Certain rocks (such as lapis lazuli, opal, and obsidian) and occasionally organic materials that are not minerals (such as amber, jet, and pearl) may also be used for jewelry and are therefore often considered to be gemstones as well. Most gemstones are hard, but some softer minerals such as brazilianite may be used in jewelry because of their color or luster or other physical properties that have aesthetic value. However, generally speaking, soft minerals are not typically used as gemstones by virtue of their brittleness and lack of durability.

Found all over the world, the industry of coloured gemstones (i.e. anything other than diamonds) is currently estimated at around US$ 1.55 billion as of 2023 and is projected to steadily increase to a value of US$ 4.46 billion by 2033.

A gem expert is a gemologist, a gem maker is called a lapidarist or gemcutter; a diamond cutter is called a diamantaire.

Characteristics and classification

A collection of gemstone pebbles made by tumbling the rough stones, except the ruby and tourmaline, with abrasive grit inside a rotating barrel. The largest pebble here is 40 mm (1.6 in) long.

The traditional classification in the West, which goes back to the ancient Greeks, begins with a distinction between precious and semi-precious; similar distinctions are made in other cultures. In modern use, the precious stones are emerald, ruby, sapphire and diamond, with all other gemstones being semi-precious. This distinction reflects the rarity of the respective stones in ancient times, as well as their quality: all are translucent, with fine color in their purest forms (except for the colorless diamond), and very hard with a hardness score of 8 to 10 on the Mohs scale. Other stones are classified by their color, translucency, and hardness. The traditional distinction does not necessarily reflect modern values; for example, while garnets are relatively inexpensive, a green garnet called tsavorite can be far more valuable than a mid-quality emerald. Another traditional term for semi-precious gemstones used in art history and archaeology is hardstone. Use of the terms 'precious' and 'semi-precious' in a commercial context is, arguably, misleading in that it suggests certain stones are more valuable than others when this is not reflected in the actual market value, although it would generally be correct if referring to desirability.

In modern times gemstones are identified by gemologists, who describe gems and their characteristics using technical terminology specific to the field of gemology. The first characteristic a gemologist uses to identify a gemstone is its chemical composition. For example, diamonds are made of carbon (C) and rubies of aluminium oxide (Al
₂O
₃). Many gems are crystals which are classified by their crystal system such as cubic or trigonal or monoclinic. Another term used is habit, the form the gem is usually found in. For example, diamonds, which have a cubic crystal system, are often found as octahedrons.

Gemstones are classified into different groups, species, and varieties. For example, ruby is the red variety of the species corundum, while any other color of corundum is considered sapphire. Other examples are the emerald (green), aquamarine (blue), red beryl (red), goshenite (colorless), heliodor (yellow), and morganite (pink), which are all varieties of the mineral species beryl.

Gems are characterized in terms of their color (hue, tone and saturation), optical phenomena, luster, refractive index, birefringence, dispersion, specific gravity, hardness, cleavage, and fracture. They may exhibit pleochroism or double refraction. They may have luminescence and a distinctive absorption spectrum. Gemstones may also be classified in terms of their "water". This is a recognized grading of the gem's luster, transparency, or "brilliance". Very transparent gems are considered "first water", while "second" or "third water" gems are those of a lesser transparency. Additionally, material or flaws within a stone may be present as inclusions.

Value

Spanish emerald and gold pendant at Victoria and Albert Museum

Enamelled gold, amethyst, and pearl pendant, about 1880, Pasquale Novissimo (1844–1914), V&A Museum number M.36-1928

Gemstones have no universally accepted grading system. Diamonds are graded using a system developed by the Gemological Institute of America (GIA) in the early 1950s. Historically, all gemstones were graded using the naked eye. The GIA system included a major innovation: the introduction of 10x magnification as the standard for grading clarity. Other gemstones are still graded using the naked eye (assuming 20/20 vision).

A mnemonic device, the "four Cs" (color, cut, clarity, and carats), has been introduced to help describe the factors used to grade a diamond. With modification, these categories can be useful in understanding the grading of all gemstones. The four criteria carry different weights depending upon whether they are applied to colored gemstones or to colorless diamonds. In diamonds, the cut is the primary determinant of value, followed by clarity and color. An ideally cut diamond will sparkle, to break down light into its constituent rainbow colors (dispersion), chop it up into bright little pieces (scintillation), and deliver it to the eye (brilliance). In its rough crystalline form, a diamond will do none of these things; it requires proper fashioning and this is called "cut". In gemstones that have color, including colored diamonds, the purity, and beauty of that color is the primary determinant of quality.

Physical characteristics that make a colored stone valuable are color, clarity to a lesser extent (emeralds will always have a number of inclusions), cut, unusual optical phenomena within the stone such as color zoning (the uneven distribution of coloring within a gem) and asteria (star effects).

Apart from the more generic and commonly used gemstones such as from diamonds, rubies, sapphires, and emeralds, pearls and opal have also been defined as precious in the jewellery trade. Up to the discoveries of bulk amethyst in Brazil in the 19th century, amethyst was considered a "precious stone" as well, going back to ancient Greece. Even in the last century certain stones such as aquamarine, peridot and cat's eye (cymophane) have been popular and hence been regarded as precious, thus reinforcing the notion that a mineral's rarity may have been implicated in its classification as a precious stone and thus contribute to its value.

Today the gemstone trade no longer makes such a distinction. Many gemstones are used in even the most expensive jewelry, depending on the brand-name of the designer, fashion trends, market supply, treatments, etc. Nevertheless, diamonds, rubies, sapphires, and emeralds still have a reputation that exceeds those of other gemstones.

Rare or unusual gemstones, generally understood to include those gemstones which occur so infrequently in gem quality that they are scarcely known except to connoisseurs, include andalusite, axinite, cassiterite, clinohumite, painite and red beryl.

Gemstone pricing and value are governed by factors and characteristics in the quality of the stone. These characteristics include clarity, rarity, freedom from defects, the beauty of the stone, as well as the demand for such stones. There are different pricing influencers for both colored gemstones, and for diamonds. The pricing on colored stones is determined by market supply-and-demand, but diamonds are more intricate.

In the addition to the aesthetic and adorning/ornamental purpose of gemstones, there are many proponents of energy medicine who also value gemstones on the basis of their alleged healing powers.

A gemstone that has been rising in popularity is Cuprian Elbaite Tourmaline which is also called "Paraiba Tourmaline". It was first discovered in the late 1980s in Paraíba, Brazil and later in Mozambique and Nigeria. It is famous for its glowing neon blue color. Paraiba Tourmaline has become one of the most popular gemstones in recent times thanks to its color and is considered to be one of the important gemstones after rubies, emeralds, and sapphires according to Gübelin Gemlab. Even though it is a tourmaline, Paraiba Tourmaline is one of the most expensive gemstones.

Grading

There are a number of laboratories which grade and provide reports on gemstones.

• Gemological Institute of America (GIA), the main provider of education services and diamond grading reports
• International Gemological Institute (IGI), independent laboratory for grading and evaluation of diamonds, jewelry, and colored stones
• Hoge Raad Voor Diamant (HRD Antwerp), The Diamond High Council, Belgium is one of Europe's oldest laboratories; its main stakeholder is the Antwerp World Diamond Centre
• American Gemological Society (AGS) is not as widely recognized nor as old as the GIA
• American Gem Trade Laboratory which is part of the American Gem Trade Association (AGTA), a trade organization of jewelers and dealers of colored stones
• American Gemological Laboratories (AGL), owned by Christopher P. Smith
• European Gemological Laboratory (EGL), founded in 1974 by Guy Margel in Belgium
• Gemmological Association of All Japan (GAAJ-ZENHOKYO), Zenhokyo, Japan, active in gemological research
• The Gem and Jewelry Institute of Thailand (Public Organization) or GIT, Thailand's national institute for gemological research and gem testing, Bangkok
• Gemmology Institute of Southern Africa, Africa's premium gem laboratory
• Asian Institute of Gemological Sciences (AIGS), the oldest gemological institute in South East Asia, involved in gemological education and gem testing
• Swiss Gemmological Institute (SSEF), founded by Henry Hänni, focusing on colored gemstones and the identification of natural pearls
• Gübelin Gem Lab, the traditional Swiss lab founded by Eduard Gübelin

Each laboratory has its own methodology to evaluate gemstones. A stone can be called "pink" by one lab while another lab calls it "padparadscha". One lab can conclude a stone is untreated, while another lab might conclude that it is heat-treated. To minimize such differences, seven of the most respected labs, AGTA-GTL (New York), CISGEM (Milano), GAAJ-ZENHOKYO (Tokyo), GIA (Carlsbad), GIT (Bangkok), Gübelin (Lucerne) and SSEF (Basel), have established the Laboratory Manual Harmonisation Committee (LMHC), for the standardization of wording reports, promotion of certain analytical methods and interpretation of results. Country of origin has sometimes been difficult to determine, due to the constant discovery of new source locations. Determining a "country of origin" is thus much more difficult than determining other aspects of a gem (such as cut, clarity, etc.).

Gem dealers are aware of the differences between gem laboratories and will make use of the discrepancies to obtain the best possible certificate.

Cutting and polishing

A diamond cutter in Amsterdam

A few gemstones are used as gems in the crystal or other forms in which they are found. Most, however, are cut and polished for usage as jewelry. The two main classifications are as follows:

• Stones cut as smooth, dome-shaped stones called cabochons or simply cab. These have been a popular shape since ancient time and is more durable than faceted gems.
• Stones which are cut with a faceting machine by polishing small flat windows called facets at regular intervals at exact angles.

Stones which are opaque or semi-opaque such as opal, turquoise, variscite, etc. are commonly cut as cabochons. These gems are designed to show the stone's color, luster and other surface properties as opposed to internal reflection properties like brilliance. Grinding wheels and polishing agents are used to grind, shape, and polish the smooth dome shape of the stones.

Gems that are transparent are normally faceted, a method that shows the optical properties of the stone's interior to its best advantage by maximizing reflected light which is perceived by the viewer as sparkle. There are many commonly used shapes for faceted stones. The facets must be cut at the proper angles, which varies depending on the optical properties of the gem. If the angles are too steep or too shallow, the light will pass through and not be reflected back toward the viewer. The faceting machine is used to hold the stone onto a flat lap for cutting and polishing the flat facets. Rarely, some cutters use special curved laps to cut and polish curved facets.

Colors

Nearly 300 variations of diamond color exhibited at the Aurora display at the Natural History Museum in London

A variety of semiprecious stones in a piece of jewellery

The color of any material is due to the nature of light itself. Daylight, often called white light, is all of the colors of the spectrum combined. When light strikes a material, most of the light is absorbed while a smaller amount of a particular frequency or wavelength is reflected. The part that is reflected reaches the eye as the perceived color. A ruby appears red because it absorbs all other colors of white light while reflecting red.

A material which is mostly the same can exhibit different colors. For example, ruby and sapphire have the same primary chemical composition (both are corundum) but exhibit different colors because of impurities which absorb and reflect different wavelengths of light depending on their individual compositions. Even the same named gemstone can occur in many different colors: sapphires show different shades of blue and pink and "fancy sapphires" exhibit a whole range of other colors from yellow to orange-pink, the latter called "padparadscha sapphire".

This difference in color is based on the atomic structure of the stone. Although the different stones formally have the same chemical composition and structure, they are not exactly the same. Every now and then an atom is replaced by a completely different atom, sometimes as few as one in a million atoms. These so-called impurities are sufficient to absorb certain colors and leave the other colors unaffected. For example, beryl, which is colorless in its pure mineral form, becomes emerald with chromium impurities. If manganese is added instead of chromium, beryl becomes pink morganite. With iron, it becomes aquamarine.Some gemstone treatments make use of the fact that these impurities can be "manipulated", thus changing the color of the gem.

Treatment

Gemstones are often treated to enhance the color or clarity of the stone. In some cases, the treatment applied to the gemstone can also increase its durability. Even though natural gemstones can be transformed using the traditional method of cutting and polishing, other treatment options allow the stone's appearance to be enhanced. Depending on the type and extent of treatment, they can affect the value of the stone. Some treatments are used widely because the resulting gem is stable, while others are not accepted most commonly because the gem color is unstable and may revert to the original tone.

Early history

Before the innovation of modern-day tools, thousands of years ago, people were recorded to use a variety of techniques to treat and enhance gemstones. Some of the earliest methods of gemstone treatment date back to the Minoan Age, for example foiling, which is where metal foil is used to enhance a gemstone's colour. Other methods recorded 2000 years ago in the book Natural History by Pliny the Elder include oiling and dyeing/staining.

Heat

Heat can either improve or spoil gemstone color or clarity. The heating process has been well known to gem miners and cutters for centuries, and in many stone types heating is a common practice. Most citrine is made by heating amethyst, and partial heating with a strong gradient results in "ametrine" – a stone partly amethyst and partly citrine. Aquamarine is often heated to remove yellow tones, or to change green colors into the more desirable blue, or enhance its existing blue color to a deeper blue.

Nearly all tanzanite is heated at low temperatures to remove brown undertones and give a more desirable blue / purple color. A considerable portion of all sapphire and ruby is treated with a variety of heat treatments to improve both color and clarity.

When jewelry containing diamonds is heated for repairs, the diamond should be protected with boric acid; otherwise, the diamond, which is pure carbon, could be burned on the surface or even burned completely up. When jewelry containing sapphires or rubies is heated, those stones should not be coated with boric acid (which can etch the surface) or any other substance. They do not have to be protected from burning, like a diamond (although the stones do need to be protected from heat stress fracture by immersing the part of the jewelry with stones in the water when metal parts are heated).

Radiation

Main article: Gemstone irradiation

The irradiation process is widely practiced in jewelry industry and enabled the creation of gemstone colors that do not exist or are extremely rare in nature. However, particularly when done in a nuclear reactor, the processes can make gemstones radioactive. Health risks related to the residual radioactivity of the treated gemstones have led to government regulations in many countries.

Virtually all blue topaz, both the lighter and the darker blue shades such as "London" blue, has been irradiated to change the color from white to blue. Most green quartz (Oro Verde) are also irradiated to achieve the yellow-green color. Diamonds are mainly irradiated to become blue-green or green, although other colors are possible. When light-to-medium-yellow diamonds are treated with gamma rays they may become green; with a high-energy electron beam, blue.

Waxing/oiling

Emeralds containing natural fissures are sometimes filled with wax or oil to disguise them. This wax or oil is also colored to make the emerald appear of better color as well as clarity. Turquoise is also commonly treated in a similar manner.

Fracture filling

The foreign material inside this fracture-filled emerald appears rainbow-colored under darkfield illumination.

Fracture filling has been in use with different gemstones such as diamonds, emeralds, and sapphires. In 2006 "glass-filled rubies" received publicity. Rubies over 10 carats (2 g) with large fractures were filled with lead glass, thus dramatically improving the appearance (of larger rubies in particular). Such treatments are fairly easy to detect.

Bleaching

Pearls are a gemstone that is commonly treated with hydrogen peroxide to remove unwanted colours

Another treatment method that is commonly used to treat gemstones is bleaching. This method uses a chemical in order to reduce the colour of the gem. After bleaching, a combination treatment can be done by dying the gemstone once the unwanted colours are removed. Hydrogen peroxide is the most commonly used product used to alter gemstones and have notably been used to treat jade and pearls. The treatment of bleaching can also be followed by impregnation, which allows the gemstone's durability to be increased.

Socioeconomic issues in the gemstone industry

The socio-economic dynamics of the gemstone industry are shaped by market forces and consumer preferences and typically go undiscussed. Changes in demand and prices can significantly affect the livelihoods of those involved in gemstone mining and trade, particularly in developing countries where the industry serves as a crucial source of income.

A situation that arises as a result of this is the exploitation of natural resources and labor within gemstone mining operations. Many mines, particularly in developing countries, face challenges such as inadequate safety measures, low wages, and poor working conditions. Miners, often from disadvantaged backgrounds, endure hazardous working conditions and receive meager wages, contributing to cycles of poverty and exploitation. Gemstone mining operations are frequently conducted in remote or underdeveloped areas, lacking proper infrastructure and access to essential services such as healthcare and education. This further contributes to the pre-existing socio-economic disparities and obstructs community development such that the benefits of gemstone extraction may not adequately reach those directly involved in the process.

Another such issue revolves around environmental degradation resulting from mining activities. Environmental degradation can pose long-term threats to ecosystems and biodiversity, further worsening the socio-economic state in affected regions. Unregulated mining practices often result in deforestation, soil erosion, and water contamination thus threatening ecosystems and biodiversity. Unregulated mining activity can also cause depletion of natural resources, thus diminishing the prospects for sustainable development. The environmental impact of gemstone mining not only poses a threat to ecosystems but also undermines the long-term viability of the industry by diminishing the quality and quantity of available resources.

Furthermore, the gemstone industry is also susceptible to issues related to transparency and ethics, which impact both producers and consumers. The lack of standardized certification processes and the prevalence of illicit practices undermine market integrity and trust. The lack of transparency and accountability in the supply chain aggravates pre-existing inequalities, as middlemen and corporations often capture a disproportionate share of the profits. As a result the unequal distribution of profits along the supply chain does little to improve socio-economic inequalities, particularly in regions where gemstones are mined.

Addressing these socio-economic challenges requires intensive effort from various stakeholders, including governments, industry executives, and society, to promote sustainable practices and ensure equitable outcomes for all involved parties. Implementing and enforcing regulations to ensure fair labor practices, environmental sustainability, and ethical sourcing is essential. Additionally, investing in community development projects, such as education and healthcare initiatives, can help alleviate poverty and empower marginalized communities dependent on the gemstone industry. Collaboration across sectors is crucial for fostering a more equitable and sustainable gemstone trade that benefits both producers and consumers while respecting human rights and environmental integrity.

Synthetic and artificial gemstones

Synthetic gemstones are distinct from imitation or simulated gems.

Synthetic gems are physically, optically, and chemically identical to the natural stone, but are created in a laboratory. Imitation or simulated stones are chemically different from the natural stone, but may appear quite similar to it; they can be more easily manufactured synthetic gemstones of a different mineral (spinel), glass, plastic, resins, or other compounds.

Examples of simulated or imitation stones include cubic zirconia, composed of zirconium oxide, synthetic moissanite, and uncolored, synthetic corundum or spinels; all of which are diamond simulants. The simulants imitate the look and color of the real stone but possess neither their chemical nor physical characteristics. In general, all are less hard than diamond. Moissanite actually has a higher refractive index than diamond, and when presented beside an equivalently sized and cut diamond will show more "fire".

Cultured, synthetic, or "lab-created" gemstones are not imitations: The bulk mineral and trace coloring elements are the same in both. For example, diamonds, rubies, sapphires, and emeralds have been manufactured in labs that possess chemical and physical characteristics identical to the naturally occurring variety. Synthetic (lab created) corundum, including ruby and sapphire, is very common and costs much less than the natural stones. Small synthetic diamonds have been manufactured in large quantities as industrial abrasives, although larger gem-quality synthetic diamonds are becoming available in multiple carats.

Whether a gemstone is a natural stone or synthetic, the chemical, physical, and optical characteristics are the same: They are composed of the same mineral and are colored by the same trace materials, have the same hardness and density and strength, and show the same color spectrum, refractive index, and birefringence (if any). Lab-created stones tend to have a more vivid color since impurities common in natural stones are not present in the synthetic stone. Synthetics are made free of common naturally occurring impurities that reduce gem clarity or color unless intentionally added in order to provide a more drab, natural appearance, or to deceive an assayer. On the other hand, synthetics often show flaws not seen in natural stones, such as minute particles of corroded metal from lab trays used during synthesis.

Types

Some gemstones are more difficult to synthesize than others and not all stones are commercially viable to attempt to synthesize. These are the most common on the market currently.

Synthetic corundum

Synthetic corundum includes ruby (red variation) and sapphire (other color variations), both of which are considered highly desired and valued. Ruby was the first gemstone to be synthesized by Auguste Verneuil with his development of the flame-fusion process in 1902. Synthetic corundum continues to be made typically by flame-fusion as it is most cost-effective, but can also be produced through flux growth and hydrothermal growth.

Synthetic beryls

The most common synthesized beryl is emerald (green). Yellow, red and blue beryls are possible but much more rare. Synthetic emerald became possible with the development of the flux growth process and is produced in this way and well as hydrothermal growth.

Synthetic quartz

Types of synthetic quartz include citrine, rose quartz, and amethyst. Natural occurring quartz is not rare, but is nevertheless synthetically produced as it has practical application outside of aesthetic purposes. Quartz generates an electric current when under pressure and is used in watches, clocks, and oscillators.

Synthetic spinel

Synthetic spinel was first produced by accident. It can be created in any color making it popular to simulate various natural gemstones. It is created through flux growth and hydrothermal growth.

Creation process

There are two main categories for creation of these minerals: melt or solution processes.

Verneuil flame fusion process (melt process)

Verneuil furnace

The flame fusion process was the first process used which successfully created large quantities of synthetic gemstones to be sold on the market. This remains the most cost effective and common method of creating corundums today.

The flame fusion process is completed in a Verneuil furnace. The furnace consists of an inverted blowpipe burner which produces an extremely hot oxyhydrogen flame, a powder dispenser, and a ceramic pedestal. A chemical powder which corresponds to the desired gemstone is passed through this flames. This melts the ingredients which drop on to a plate and solidify into a crystal called a boule. For corundum the flame must be 2000 °C. This process takes hours and yields a crystal with the same properties as its natural counterpart.

To produce corundum, a pure aluminium powder is used with different additives to achieve different colors.

• Chromic oxide for ruby
• Iron and titanium oxide for blue sapphire
• Nickel oxide for yellow sapphire
• Nickel, chromium and iron for orange sapphire
• Manganese for pink sapphire
• Copper for blue-green sapphire
• Cobalt for dark blue sapphire

Czochralski process (melt process)

In 1918 this process was developed by J. Czocharalski and is also referred to as the "crystal pulling" method. In this process, the required gemstone materials are added to a crucible. A seed stone is placed into the melt in the crucible. As the gem begins to crystallize on the seed, the seed is pulled away and the gem continues to grow. This is used for corundum but is currently the least popular method.

Flux growth (solution process)

The flux growth process was the first process able to synthesize emerald. Flux growth begins with a crucible which can withstand high heat; either graphite or platinum which is filled with a molten liquid referred to as flux. The specific gem ingredients are added and dissolved in this fluid and recrystallize to form the desired gemstone.This is a longer process compared to the flame fusion process and can take two months up to a year depending on the desired final size.

Hydrothermal growth (solution process)

The hydrothermal growth process attempts to imitate the natural growth process of minerals. The required gem materials are sealed in a container of water and placed under extreme pressure. The water is heated beyond its boiling point which allows normally insoluble materials to dissolve. As more material cannot be added once the container is sealed, in order to create a larger gem the process would begin with a "seed" stone from a previous batch which the new material will crystallize on. This process takes a few weeks to complete.

Characteristics

Synthetic gemstones share chemical and physical properties with natural gemstones, but there are some slight differences that can be used to discern synthetic from natural. These differences are slight and often require microscopy as a tool to distinguish differences. Undetectable synthetics pose a threat to the market if they are able to be sold as rare natural gemstones. Because of this there are certain characteristic gemologists look for. Each crystal is characteristic to the environment and growth process under which it was created.

Visible banding in an apatite gemstone

Gemstones created from the flame-fusion process may have

• small air bubbles which were trapped inside the boule during formation process
• visible banding from formation of the boule
• chatter marks which on the surface which appear crack like which are caused from damage during polishing of the gemstone

Gemstones created from flux melt process may have

• small cavities which are filled with flux solution
• inclusions in the gemstone from crucible used

Gemstones created from hydrothermal growth may have

• inclusions from container used

History

Auguste Verneuil – creator of flame-fusion process 1902

Prior to development of synthesising processes the alternatives on the market to natural gemstones were imitations or fake. In 1837, the first successful synthesis of ruby occurred. French chemist Marc Gaudin managed to produce small crystals of ruby from melting together potassium aluminium sulphate and potassium chromate through what would later be known as the flux melt process. Following this, another French chemist Fremy was able to grow large quantities of small ruby crystals using a lead flux.

A few years later an alternative to flux melt was developed which led to the introduction of what was labeled "reconstructed ruby" to the market. Reconstructed ruby was sold as a process which produced larger rubies from melting together bits of natural ruby. In later attempts to recreate this process it was found to not be possible and is believed reconstructed rubies were most likely created using a multi-step method of melting of ruby powder.

Auguste Verneuil, a student of Fremy, went on to develop flame-fusion as an alternative to the flux-melt method. He developed large furnaces which were able to produce large quantities of corundums more efficiently and shifted the gemstone market dramatically. This process is still used today and the furnaces have not changed much from the original design. World production of corundum using this method reaches 1000 million carats a year.

List of rare gemstones

• Painite was discovered in 1956 in Ohngaing in Myanmar. The mineral was named in honor of the British gemologist Arthur Charles Davy Pain. At one point it was considered the rarest mineral on Earth.
• Tanzanite was discovered in 1967 in Northern Tanzania. With its supply possibly declining in the next 30 years, this gemstone is considered to be more rare than a diamond. This type of gemstone receives its vibrant blue from being heated.
• Hibonite was discovered in 1956 in Madagascar. It was named after the discoverer, French geologist Paul Hibon. Gem quality hibonite has been found only in Myanmar.

Red Beryl - discovered in 1940

• Red beryl or bixbite was discovered in an area near Beaver, Utah in 1904 and named after the American mineralogist Maynard Bixby.
• Jeremejevite was discovered in 1883 in Russia and named after its discoverer, Pawel Wladimirowich Jeremejew (1830–1899).
• Chambersite was discovered in 1957 in Chambers County, Texas, US, and named after the deposit's location.
• Taaffeite was discovered in 1945. It was named after the discoverer, the Irish gemologist Count Edward Charles Richard Taaffe.
• Musgravite was discovered in 1967 in the Musgrave Mountains in South Australia and named for the location.

Black Opal – the rarest type of opal

• Black opal is directly mined in New South Wales, Australia, making it the rarest type of opal. Having a darker composition, this gemstone can be in a variety of colours.
• Grandidierite was discovered by Antoine François Alfred Lacroix (1863–1948) in 1902 in Tuléar Province, Madagascar. It was named in honor of the French naturalist and explorer Alfred Grandidier (1836–1912).
• Poudretteite was discovered in 1965 at the Poudrette Quarry in Canada and named after the quarry's owners and operators, the Poudrette family.
• Serendibite was discovered in Sri Lanka by Sunil Palitha Gunasekera in 1902 and named after Serendib, the old Arabic name for Sri Lanka.
• Zektzerite was discovered by Bart Cannon in 1968 on Kangaroo Ridge near Washington Pass in Okanogan County, Washington, USA. The mineral was named in honor of mathematician and geologist Jack Zektzer, who presented the material for study in 1976.

In popular culture

French singer-songwriter Nolwenn Leroy was inspired by the gemstones for her 2017 album Gemme (meaning gemstone in French) and the single of the same name.

Electron configuration

From Wikipedia, the free encyclopedia
https://en.wikipedia.org/wiki/Electron_configuration

Electron atomic and molecular orbitals

A Bohr diagram of lithium

In atomic physics and quantum chemistry, the electron configuration is the distribution of electrons of an atom or molecule (or other physical structure) in atomic or molecular orbitals. For example, the electron configuration of the neon atom is 1s² 2s² 2p⁶, meaning that the 1s, 2s, and 2p subshells are occupied by two, two, and six electrons, respectively.

Electronic configurations describe each electron as moving independently in an orbital, in an average field created by the nuclei and all the other electrons. Mathematically, configurations are described by Slater determinants or configuration state functions.

According to the laws of quantum mechanics, a level of energy is associated with each electron configuration. In certain conditions, electrons are able to move from one configuration to another by the emission or absorption of a quantum of energy, in the form of a photon.

Knowledge of the electron configuration of different atoms is useful in understanding the structure of the periodic table of elements, for describing the chemical bonds that hold atoms together, and in understanding the chemical formulas of compounds and the geometries of molecules. In bulk materials, this same idea helps explain the peculiar properties of lasers and semiconductors.

Shells and subshells

Main article: Electron shell

	s (l = 0)	p (l = 1)
	m = 0	m = 0	m = ±1
	s	pz	px	py
n = 1
n = 2

Electron configuration was first conceived under the Bohr model of the atom, and it is still common to speak of shells and subshells despite the advances in understanding of the quantum-mechanical nature of electrons.

An electron shell is the set of allowed states that share the same principal quantum number, n, that electrons may occupy. In each term of an electron configuration, n is the positive integer that precedes each orbital letter (helium's electron configuration is 1s², therefore n = 1, and the orbital contains two electrons). An atom's nth electron shell can accommodate 2n² electrons. For example, the first shell can accommodate two electrons, the second shell eight electrons, the third shell eighteen, and so on. The factor of two arises because the number of allowed states doubles with each successive shell due to electron spin—each atomic orbital admits up to two otherwise identical electrons with opposite spin, one with a spin +1⁄2 (usually denoted by an up-arrow) and one with a spin of −1⁄2 (with a down-arrow).

A subshell is the set of states defined by a common azimuthal quantum number, l, within a shell. The value of l is in the range from 0 to n − 1. The values l = 0, 1, 2, 3 correspond to the s, p, d, and f labels, respectively. For example, the 3d subshell has n = 3 and l = 2. The maximum number of electrons that can be placed in a subshell is given by 2(2l + 1). This gives two electrons in an s subshell, six electrons in a p subshell, ten electrons in a d subshell and fourteen electrons in an f subshell.

The numbers of electrons that can occupy each shell and each subshell arise from the equations of quantum mechanics, in particular the Pauli exclusion principle, which states that no two electrons in the same atom can have the same values of the four quantum numbers.

Notation

See also: Atomic orbital

Physicists and chemists use a standard notation to indicate the electron configurations of atoms and molecules. For atoms, the notation consists of a sequence of atomic subshell labels (e.g. for phosphorus the sequence 1s, 2s, 2p, 3s, 3p) with the number of electrons assigned to each subshell placed as a superscript. For example, hydrogen has one electron in the s-orbital of the first shell, so its configuration is written 1s¹. Lithium has two electrons in the 1s-subshell and one in the (higher-energy) 2s-subshell, so its configuration is written 1s² 2s¹ (pronounced "one-s-two, two-s-one"). Phosphorus (atomic number 15) is as follows: 1s² 2s² 2p⁶ 3s² 3p³.

For atoms with many electrons, this notation can become lengthy and so an abbreviated notation is used. The electron configuration can be visualized as the core electrons, equivalent to the noble gas of the preceding period, and the valence electrons: each element in a period differs only by the last few subshells. Phosphorus, for instance, is in the third period. It differs from the second-period neon, whose configuration is 1s² 2s² 2p⁶, only by the presence of a third shell. The portion of its configuration that is equivalent to neon is abbreviated as [Ne], allowing the configuration of phosphorus to be written as [Ne] 3s² 3p³ rather than writing out the details of the configuration of neon explicitly. This convention is useful as it is the electrons in the outermost shell that most determine the chemistry of the element.

For a given configuration, the order of writing the orbitals is not completely fixed since only the orbital occupancies have physical significance. For example, the electron configuration of the titanium ground state can be written as either [Ar] 4s² 3d² or [Ar] 3d² 4s². The first notation follows the order based on the Madelung rule for the configurations of neutral atoms; 4s is filled before 3d in the sequence Ar, K, Ca, Sc, Ti. The second notation groups all orbitals with the same value of n together, corresponding to the "spectroscopic" order of orbital energies that is the reverse of the order in which electrons are removed from a given atom to form positive ions; 3d is filled before 4s in the sequence Ti⁴⁺, Ti³⁺, Ti²⁺, Ti⁺, Ti.

The superscript 1 for a singly occupied subshell is not compulsory; for example aluminium may be written as either [Ne] 3s² 3p¹ or [Ne] 3s² 3p. In atoms where a subshell is unoccupied despite higher subshells being occupied (as is the case in some ions, as well as certain neutral atoms shown to deviate from the Madelung rule), the empty subshell is either denoted with a superscript 0 or left out altogether. For example, neutral palladium may be written as either [Kr] 4d¹⁰ 5s⁰ or simply [Kr] 4d¹⁰, and the lanthanum(III) ion may be written as either [Xe] 4f⁰ or simply [Xe].

It is quite common to see the letters of the orbital labels (s, p, d, f) written in an italic or slanting typeface, although the International Union of Pure and Applied Chemistry (IUPAC) recommends a normal typeface (as used here). The choice of letters originates from a now-obsolete system of categorizing spectral lines as "sharp", "principal", "diffuse" and "fundamental" (or "fine"), based on their observed fine structure: their modern usage indicates orbitals with an azimuthal quantum number, l, of 0, 1, 2 or 3 respectively. After f, the sequence continues alphabetically g, h, i... (l = 4, 5, 6...), skipping j, although orbitals of these types are rarely required.

The electron configurations of molecules are written in a similar way, except that molecular orbital labels are used instead of atomic orbital labels (see below).

Energy of ground state and excited states

The energy associated to an electron is that of its orbital. The energy of a configuration is often approximated as the sum of the energy of each electron, neglecting the electron-electron interactions. The configuration that corresponds to the lowest electronic energy is called the ground state. Any other configuration is an excited state.

As an example, the ground state configuration of the sodium atom is 1s² 2s² 2p⁶ 3s¹, as deduced from the Aufbau principle (see below). The first excited state is obtained by promoting a 3s electron to the 3p subshell, to obtain the 1s² 2s² 2p⁶ 3p¹ configuration, abbreviated as the 3p level. Atoms can move from one configuration to another by absorbing or emitting energy. In a sodium-vapor lamp for example, sodium atoms are excited to the 3p level by an electrical discharge, and return to the ground state by emitting yellow light of wavelength 589 nm.

Usually, the excitation of valence electrons (such as 3s for sodium) involves energies corresponding to photons of visible or ultraviolet light. The excitation of core electrons is possible, but requires much higher energies, generally corresponding to X-ray photons. This would be the case for example to excite a 2p electron of sodium to the 3s level and form the excited 1s² 2s² 2p⁵ 3s² configuration.

The remainder of this article deals only with the ground-state configuration, often referred to as "the" configuration of an atom or molecule.

History

Irving Langmuir was the first to propose in his 1919 article "The Arrangement of Electrons in Atoms and Molecules" in which, building on Gilbert N. Lewis's cubical atom theory and Walther Kossel's chemical bonding theory, he outlined his "concentric theory of atomic structure". Langmuir had developed his work on electron atomic structure from other chemists as is shown in the development of the History of the periodic table and the Octet rule.

Niels Bohr (1923) incorporated Langmuir's model that the periodicity in the properties of the elements might be explained by the electronic structure of the atom. His proposals were based on the then current Bohr model of the atom, in which the electron shells were orbits at a fixed distance from the nucleus. Bohr's original configurations would seem strange to a present-day chemist: sulfur was given as 2.4.4.6 instead of 1s² 2s² 2p⁶ 3s² 3p⁴ (2.8.6). Bohr used 4 and 6 following Alfred Werner's 1893 paper. In fact, the chemists accepted the concept of atoms long before the physicists. Langmuir began his paper referenced above by saying,

«…The problem of the structure of atoms has been attacked mainly by physicists who have given little consideration to the chemical properties which must ultimately be explained by a theory of atomic structure. The vast store of knowledge of chemical properties and relationships, such as is summarized by the Periodic Table, should serve as a better foundation for a theory of atomic structure than the relatively meager experimental data along purely physical lines... These electrons arrange themselves in a series of concentric shells, the first shell containing two electrons, while all other shells tend to hold eight.…»

The valence electrons in the atom were described by Richard Abegg in 1904.

In 1924, E. C. Stoner incorporated Sommerfeld's third quantum number into the description of electron shells, and correctly predicted the shell structure of sulfur to be 2.8.6. However neither Bohr's system nor Stoner's could correctly describe the changes in atomic spectra in a magnetic field (the Zeeman effect).

Bohr was well aware of this shortcoming (and others), and had written to his friend Wolfgang Pauli in 1923 to ask for his help in saving quantum theory (the system now known as "old quantum theory"). Pauli hypothesized successfully that the Zeeman effect can be explained as depending only on the response of the outermost (i.e., valence) electrons of the atom. Pauli was able to reproduce Stoner's shell structure, but with the correct structure of subshells, by his inclusion of a fourth quantum number and his exclusion principle (1925):

It should be forbidden for more than one electron with the same value of the main quantum number n to have the same value for the other three quantum numbers k [l], j [m_l] and m [m_s].

The Schrödinger equation, published in 1926, gave three of the four quantum numbers as a direct consequence of its solution for the hydrogen atom: this solution yields the atomic orbitals that are shown today in textbooks of chemistry (and above). The examination of atomic spectra allowed the electron configurations of atoms to be determined experimentally, and led to an empirical rule (known as Madelung's rule (1936), see below) for the order in which atomic orbitals are filled with electrons.

Atoms: Aufbau principle and Madelung rule

See also: Electron configurations of the elements (data page)

The aufbau principle (from the German Aufbau, "building up, construction") was an important part of Bohr's original concept of electron configuration. It may be stated as:

a maximum of two electrons are put into orbitals in the order of increasing orbital energy: the lowest-energy subshells are filled before electrons are placed in higher-energy orbitals.

The approximate order of filling of atomic orbitals, following the arrows from 1s to 7p. (After 7p the order includes subshells outside the range of the diagram, starting with 8s.)

The principle works very well (for the ground states of the atoms) for the known 118 elements, although it is sometimes slightly wrong. The modern form of the aufbau principle describes an order of orbital energies given by Madelung's rule (or Klechkowski's rule). This rule was first stated by Charles Janet in 1929, rediscovered by Erwin Madelung in 1936, and later given a theoretical justification by V. M. Klechkowski:

1. Subshells are filled in the order of increasing n + l.
2. Where two subshells have the same value of n + l, they are filled in order of increasing n.

This gives the following order for filling the orbitals:

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, (8s, 5g, 6f, 7d, 8p, and 9s)

In this list the subshells in parentheses are not occupied in the ground state of the heaviest atom now known (Og, Z = 118).

The aufbau principle can be applied, in a modified form, to the protons and neutrons in the atomic nucleus, as in the shell model of nuclear physics and nuclear chemistry.

Periodic table

Main article: Block (periodic table)

Electron configuration table showing blocks.

The form of the periodic table is closely related to the atomic electron configuration for each element. For example, all the elements of group 2 (the table's second column) have an electron configuration of [E] ns² (where [E] is a noble gas configuration), and have notable similarities in their chemical properties. The periodicity of the periodic table in terms of periodic table blocks is due to the number of electrons (2, 6, 10, and 14) needed to fill s, p, d, and f subshells. These blocks appear as the rectangular sections of the periodic table. The single exception is helium, which despite being an s-block atom is conventionally placed with the other noble gasses in the p-block due to its chemical inertness, a consequence of its full outer shell (though there is discussion in the contemporary literature on whether this exception should be retained).

The electrons in the valence (outermost) shell largely determine each element's chemical properties. The similarities in the chemical properties were remarked on more than a century before the idea of electron configuration.

Shortcomings of the aufbau principle

The aufbau principle rests on a fundamental postulate that the order of orbital energies is fixed, both for a given element and between different elements; in both cases this is only approximately true. It considers atomic orbitals as "boxes" of fixed energy into which can be placed two electrons and no more. However, the energy of an electron "in" an atomic orbital depends on the energies of all the other electrons of the atom (or ion, or molecule, etc.). There are no "one-electron solutions" for systems of more than one electron, only a set of many-electron solutions that cannot be calculated exactly (although there are mathematical approximations available, such as the Hartree–Fock method).

The fact that the aufbau principle is based on an approximation can be seen from the fact that there is an almost-fixed filling order at all, that, within a given shell, the s-orbital is always filled before the p-orbitals. In a hydrogen-like atom, which only has one electron, the s-orbital and the p-orbitals of the same shell have exactly the same energy, to a very good approximation in the absence of external electromagnetic fields. (However, in a real hydrogen atom, the energy levels are slightly split by the magnetic field of the nucleus, and by the quantum electrodynamic effects of the Lamb shift.)

Ionization of the transition metals

The naïve application of the aufbau principle leads to a well-known paradox (or apparent paradox) in the basic chemistry of the transition metals. Potassium and calcium appear in the periodic table before the transition metals, and have electron configurations [Ar] 4s¹ and [Ar] 4s² respectively, i.e. the 4s-orbital is filled before the 3d-orbital. This is in line with Madelung's rule, as the 4s-orbital has n + l = 4 (n = 4, l = 0) while the 3d-orbital has n + l = 5 (n = 3, l = 2). After calcium, most neutral atoms in the first series of transition metals (scandium through zinc) have configurations with two 4s electrons, but there are two exceptions. Chromium and copper have electron configurations [Ar] 3d⁵ 4s¹ and [Ar] 3d¹⁰ 4s¹ respectively, i.e. one electron has passed from the 4s-orbital to a 3d-orbital to generate a half-filled or filled subshell. In this case, the usual explanation is that "half-filled or completely filled subshells are particularly stable arrangements of electrons". However, this is not supported by the facts, as tungsten (W) has a Madelung-following d⁴ s² configuration and not d⁵ s¹, and niobium (Nb) has an anomalous d⁴ s¹ configuration that does not give it a half-filled or completely filled subshell.

The apparent paradox arises when electrons are removed from the transition metal atoms to form ions. The first electrons to be ionized come not from the 3d-orbital, as one would expect if it were "higher in energy", but from the 4s-orbital. This interchange of electrons between 4s and 3d is found for all atoms of the first series of transition metals. The configurations of the neutral atoms (K, Ca, Sc, Ti, V, Cr, ...) usually follow the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, ...; however the successive stages of ionization of a given atom (such as Fe⁴⁺, Fe³⁺, Fe²⁺, Fe⁺, Fe) usually follow the order 1s, 2s, 2p, 3s, 3p, 3d, 4s, ...

This phenomenon is only paradoxical if it is assumed that the energy order of atomic orbitals is fixed and unaffected by the nuclear charge or by the presence of electrons in other orbitals. If that were the case, the 3d-orbital would have the same energy as the 3p-orbital, as it does in hydrogen, yet it clearly does not. There is no special reason why the Fe²⁺ ion should have the same electron configuration as the chromium atom, given that iron has two more protons in its nucleus than chromium, and that the chemistry of the two species is very different. Melrose and Eric Scerri have analyzed the changes of orbital energy with orbital occupations in terms of the two-electron repulsion integrals of the Hartree–Fock method of atomic structure calculation. More recently Scerri has argued that contrary to what is stated in the vast majority of sources including the title of his previous article on the subject, 3d orbitals rather than 4s are in fact preferentially occupied.

In chemical environments, configurations can change even more: Th³⁺ as a bare ion has a configuration of [Rn] 5f¹, yet in most Th^III compounds the thorium atom has a 6d¹ configuration instead. Mostly, what is present is rather a superposition of various configurations. For instance, copper metal is poorly described by either an [Ar] 3d¹⁰ 4s¹ or an [Ar] 3d⁹ 4s² configuration, but is rather well described as a 90% contribution of the first and a 10% contribution of the second. Indeed, visible light is already enough to excite electrons in most transition metals, and they often continuously "flow" through different configurations when that happens (copper and its group are an exception).

Similar ion-like 3d^x 4s⁰ configurations occur in transition metal complexes as described by the simple crystal field theory, even if the metal has oxidation state 0. For example, chromium hexacarbonyl can be described as a chromium atom (not ion) surrounded by six carbon monoxide ligands. The electron configuration of the central chromium atom is described as 3d⁶ with the six electrons filling the three lower-energy d orbitals between the ligands. The other two d orbitals are at higher energy due to the crystal field of the ligands. This picture is consistent with the experimental fact that the complex is diamagnetic, meaning that it has no unpaired electrons. However, in a more accurate description using molecular orbital theory, the d-like orbitals occupied by the six electrons are no longer identical with the d orbitals of the free atom.

Other exceptions to Madelung's rule

There are several more exceptions to Madelung's rule among the heavier elements, and as atomic number increases it becomes more and more difficult to find simple explanations such as the stability of half-filled subshells. It is possible to predict most of the exceptions by Hartree–Fock calculations, which are an approximate method for taking account of the effect of the other electrons on orbital energies. Qualitatively, for example, the 4d elements have the greatest concentration of Madelung anomalies, because the 4d–5s gap is larger than the 3d–4s and 5d–6s gaps.

For the heavier elements, it is also necessary to take account of the effects of special relativity on the energies of the atomic orbitals, as the inner-shell electrons are moving at speeds approaching the speed of light. In general, these relativistic effects tend to decrease the energy of the s-orbitals in relation to the other atomic orbitals. This is the reason why the 6d elements are predicted to have no Madelung anomalies apart from lawrencium (for which relativistic effects stabilise the p_1/2 orbital as well and cause its occupancy in the ground state), as relativity intervenes to make the 7s orbitals lower in energy than the 6d ones.

The table below shows the configurations of the f-block (green) and d-block (blue) atoms. It shows the ground state configuration in terms of orbital occupancy, but it does not show the ground state in terms of the sequence of orbital energies as determined spectroscopically. For example, in the transition metals, the 4s orbital is of a higher energy than the 3d orbitals; and in the lanthanides, the 6s is higher than the 4f and 5d. The ground states can be seen in the Electron configurations of the elements (data page). However this also depends on the charge: a calcium atom has 4s lower in energy than 3d, but a Ca²⁺ cation has 3d lower in energy than 4s. In practice the configurations predicted by the Madelung rule are at least close to the ground state even in these anomalous cases. The empty f orbitals in lanthanum, actinium, and thorium contribute to chemical bonding, as do the empty p orbitals in transition metals.

Vacant s, d, and f orbitals have been shown explicitly, as is occasionally done, to emphasise the filling order and to clarify that even orbitals unoccupied in the ground state (e.g. lanthanum 4f or palladium 5s) may be occupied and bonding in chemical compounds. (The same is also true for the p-orbitals, which are not explicitly shown because they are only actually occupied for lawrencium in gas-phase ground states.)

Electron shells filled in violation of Madelung's rule (red)
Predictions for elements 109–112

Period 4				Period 5				Period 6				Period 7
Element	Z	Electron Configuration		Element	Z	Electron Configuration		Element	Z	Electron Configuration		Element	Z	Electron Configuration
								Lanthanum	57	[Xe] 6s2 4f0 5d1		Actinium	89	[Rn] 7s2 5f0 6d1
								Cerium	58	[Xe] 6s2 4f1 5d1		Thorium	90	[Rn] 7s2 5f0 6d2
								Praseodymium	59	[Xe] 6s2 4f3 5d0		Protactinium	91	[Rn] 7s2 5f2 6d1
								Neodymium	60	[Xe] 6s2 4f4 5d0		Uranium	92	[Rn] 7s2 5f3 6d1
								Promethium	61	[Xe] 6s2 4f5 5d0		Neptunium	93	[Rn] 7s2 5f4 6d1
								Samarium	62	[Xe] 6s2 4f6 5d0		Plutonium	94	[Rn] 7s2 5f6 6d0
								Europium	63	[Xe] 6s2 4f7 5d0		Americium	95	[Rn] 7s2 5f7 6d0
								Gadolinium	64	[Xe] 6s2 4f7 5d1		Curium	96	[Rn] 7s2 5f7 6d1
								Terbium	65	[Xe] 6s2 4f9 5d0		Berkelium	97	[Rn] 7s2 5f9 6d0
								Dysprosium	66	[Xe] 6s2 4f10 5d0		Californium	98	[Rn] 7s2 5f10 6d0
								Holmium	67	[Xe] 6s2 4f11 5d0		Einsteinium	99	[Rn] 7s2 5f11 6d0
								Erbium	68	[Xe] 6s2 4f12 5d0		Fermium	100	[Rn] 7s2 5f12 6d0
								Thulium	69	[Xe] 6s2 4f13 5d0		Mendelevium	101	[Rn] 7s2 5f13 6d0
								Ytterbium	70	[Xe] 6s2 4f14 5d0		Nobelium	102	[Rn] 7s2 5f14 6d0
Scandium	21	[Ar] 4s2 3d1		Yttrium	39	[Kr] 5s2 4d1		Lutetium	71	[Xe] 6s2 4f14 5d1		Lawrencium	103	[Rn] 7s2 5f14 6d0 7p1
Titanium	22	[Ar] 4s2 3d2		Zirconium	40	[Kr] 5s2 4d2		Hafnium	72	[Xe] 6s2 4f14 5d2		Rutherfordium	104	[Rn] 7s2 5f14 6d2
Vanadium	23	[Ar] 4s2 3d3		Niobium	41	[Kr] 5s1 4d4		Tantalum	73	[Xe] 6s2 4f14 5d3		Dubnium	105	[Rn] 7s2 5f14 6d3
Chromium	24	[Ar] 4s1 3d5		Molybdenum	42	[Kr] 5s1 4d5		Tungsten	74	[Xe] 6s2 4f14 5d4		Seaborgium	106	[Rn] 7s2 5f14 6d4
Manganese	25	[Ar] 4s2 3d5		Technetium	43	[Kr] 5s2 4d5		Rhenium	75	[Xe] 6s2 4f14 5d5		Bohrium	107	[Rn] 7s2 5f14 6d5
Iron	26	[Ar] 4s2 3d6		Ruthenium	44	[Kr] 5s1 4d7		Osmium	76	[Xe] 6s2 4f14 5d6		Hassium	108	[Rn] 7s2 5f14 6d6
Cobalt	27	[Ar] 4s2 3d7		Rhodium	45	[Kr] 5s1 4d8		Iridium	77	[Xe] 6s2 4f14 5d7		Meitnerium	109	[Rn] 7s2 5f14 6d7
Nickel	28	[Ar] 4s2 3d8 or [Ar] 4s1 3d9		Palladium	46	[Kr] 5s0 4d10		Platinum	78	[Xe] 6s1 4f14 5d9		Darmstadtium	110	[Rn] 7s2 5f14 6d8
Copper	29	[Ar] 4s1 3d10		Silver	47	[Kr] 5s1 4d10		Gold	79	[Xe] 6s1 4f14 5d10		Roentgenium	111	[Rn] 7s2 5f14 6d9
Zinc	30	[Ar] 4s2 3d10		Cadmium	48	[Kr] 5s2 4d10		Mercury	80	[Xe] 6s2 4f14 5d10		Copernicium	112	[Rn] 7s2 5f14 6d10

The various anomalies describe the free atoms and do not necessarily predict chemical behavior. Thus for example neodymium typically forms the +3 oxidation state, despite its configuration [Xe] 4f⁴ 5d⁰ 6s² that if interpreted naïvely would suggest a more stable +2 oxidation state corresponding to losing only the 6s electrons. Contrariwise, uranium as [Rn] 5f³ 6d¹ 7s² is not very stable in the +3 oxidation state either, preferring +4 and +6.

The electron-shell configuration of elements beyond hassium has not yet been empirically verified, but they are expected to follow Madelung's rule without exceptions until element 120. Element 121 should have the anomalous configuration [Og] 8s² 5g⁰ 6f⁰ 7d⁰ 8p¹, having a p rather than a g electron. Electron configurations beyond this are tentative and predictions differ between models, but Madelung's rule is expected to break down due to the closeness in energy of the 5g, 6f, 7d, and 8p_1/2 orbitals. That said, the filling sequence 8s, 5g, 6f, 7d, 8p is predicted to hold approximately, with perturbations due to the huge spin-orbit splitting of the 8p and 9p shells, and the huge relativistic stabilisation of the 9s shell.

Open and closed shells

In the context of atomic orbitals, an open shell is a valence shell which is not completely filled with electrons or that has not given all of its valence electrons through chemical bonds with other atoms or molecules during a chemical reaction. Conversely a closed shell is obtained with a completely filled valence shell. This configuration is very stable.

For molecules, "open shell" signifies that there are unpaired electrons. In molecular orbital theory, this leads to molecular orbitals that are singly occupied. In computational chemistry implementations of molecular orbital theory, open-shell molecules have to be handled by either the restricted open-shell Hartree–Fock method or the unrestricted Hartree–Fock method. Conversely a closed-shell configuration corresponds to a state where all molecular orbitals are either doubly occupied or empty (a singlet state). Open shell molecules are more difficult to study computationally.

Noble gas configuration

Further information: Noble gas

Noble gas configuration is the electron configuration of noble gases. The basis of all chemical reactions is the tendency of chemical elements to acquire stability. Main-group atoms generally obey the octet rule, while transition metals generally obey the 18-electron rule. The noble gases (He, Ne, Ar, Kr, Xe, Rn) are less reactive than other elements because they already have a noble gas configuration. Oganesson is predicted to be more reactive due to relativistic effects for heavy atoms.

Period	Element	Configuration
1	He	1s2
2	Ne	1s2	2s2 2p6
3	Ar	1s2	2s2 2p6	3s2 3p6
4	Kr	1s2	2s2 2p6	3s2 3p6	4s2 3d10 4p6
5	Xe	1s2	2s2 2p6	3s2 3p6	4s2 3d10 4p6	5s2 4d10 5p6
6	Rn	1s2	2s2 2p6	3s2 3p6	4s2 3d10 4p6	5s2 4d10 5p6	6s2 4f14 5d10 6p6
7	Og	1s2	2s2 2p6	3s2 3p6	4s2 3d10 4p6	5s2 4d10 5p6	6s2 4f14 5d10 6p6	7s2 5f14 6d10 7p6

Every system has the tendency to acquire the state of stability or a state of minimum energy, and so chemical elements take part in chemical reactions to acquire a stable electronic configuration similar to that of its nearest noble gas. An example of this tendency is two hydrogen (H) atoms reacting with one oxygen (O) atom to form water (H₂O). Neutral atomic hydrogen has one electron in its valence shell, and on formation of water it acquires a share of a second electron coming from oxygen, so that its configuration is similar to that of its nearest noble gas helium (He) with two electrons in its valence shell. Similarly, neutral atomic oxygen has six electrons in its valence shell, and acquires a share of two electrons from the two hydrogen atoms, so that its configuration is similar to that of its nearest noble gas neon with eight electrons in its valence shell.

Electron configuration in molecules

Electron configuration in molecules is more complex than the electron configuration of atoms, as each molecule has a different orbital structure. The molecular orbitals are labelled according to their symmetry, rather than the atomic orbital labels used for atoms and monatomic ions; hence, the electron configuration of the dioxygen molecule, O₂, is written 1σ_g² 1σ_u² 2σ_g² 2σ_u² 3σ_g² 1π_u⁴ 1π_g², or equivalently 1σ_g² 1σ_u² 2σ_g² 2σ_u² 1π_u⁴ 3σ_g² 1π_g². The term 1π_g² represents the two electrons in the two degenerate π*-orbitals (antibonding). From Hund's rules, these electrons have parallel spins in the ground state, and so dioxygen has a net magnetic moment (it is paramagnetic). The explanation of the paramagnetism of dioxygen was a major success for molecular orbital theory.

The electronic configuration of polyatomic molecules can change without absorption or emission of a photon through vibronic couplings.

Electron configuration in solids

In a solid, the electron states become very numerous. They cease to be discrete, and effectively blend into continuous ranges of possible states (an electron band). The notion of electron configuration ceases to be relevant, and yields to band theory.

Applications

The most widespread application of electron configurations is in the rationalization of chemical properties, in both inorganic and organic chemistry. In effect, electron configurations, along with some simplified forms of molecular orbital theory, have become the modern equivalent of the valence concept, describing the number and type of chemical bonds that an atom can be expected to form.

This approach is taken further in computational chemistry, which typically attempts to make quantitative estimates of chemical properties. For many years, most such calculations relied upon the "linear combination of atomic orbitals" (LCAO) approximation, using an ever-larger and more complex basis set of atomic orbitals as the starting point. The last step in such a calculation is the assignment of electrons among the molecular orbitals according to the aufbau principle. Not all methods in computational chemistry rely on electron configuration: density functional theory (DFT) is an important example of a method that discards the model.

For atoms or molecules with more than one electron, the motion of electrons are correlated and such a picture is no longer exact. A very large number of electronic configurations are needed to exactly describe any multi-electron system, and no energy can be associated with one single configuration. However, the electronic wave function is usually dominated by a very small number of configurations and therefore the notion of electronic configuration remains essential for multi-electron systems.

A fundamental application of electron configurations is in the interpretation of atomic spectra. In this case, it is necessary to supplement the electron configuration with one or more term symbols, which describe the different energy levels available to an atom. Term symbols can be calculated for any electron configuration, not just the ground-state configuration listed in tables, although not all the energy levels are observed in practice. It is through the analysis of atomic spectra that the ground-state electron configurations of the elements were experimentally determined.