Transition metal

In chemistry, the term transition metal (or transition element) has three possible meanings:

• The IUPAC definition[1] defines a transition metal as "an element whose atom has a partially filled d sub-shell, or which can give rise to cations with an incomplete d sub-shell".
• Many scientists describe a "transition metal" as any element in the d-block of the periodic table, which includes groups 3 to 12 on the periodic table.[2][3] In actual practice, the f-block lanthanide and actinide series are also considered transition metals and are called "inner transition metals".
• Cotton and Wilkinson[4] expand the brief IUPAC definition (see above) by specifying which elements are included. As well as the elements of groups 4 to 11, they add scandium and yttrium in group 3, which have a partially filled d subshell in the metallic state. Lanthanum and actinium in group 3 are, however, classified as lanthanides and actinides respectively.

English chemist Charles Bury (1890–1968) first used the word transition in this context in 1921, when he referred to a transition series of elements during the change of an inner layer of electrons (for example n = 3 in the 4th row of the periodic table) from a stable group of 8 to one of 18, or from 18 to 32.[5][6][7] These elements are now known as the d-block.

Classification

In the d-block the atoms of the elements have between 1 and 10 d electrons.

Transition metals in the d-block
Group 3 4 5 6 7 8 9 10 11 12
Period 4 21Sc 22Ti 23V 24Cr 25Mn 26Fe 27Co 28Ni 29Cu 30Zn
5 39Y 40Zr 41Nb 42Mo 43Tc 44Ru 45Rh 46Pd 47Ag 48Cd
6 57La 72Hf 73Ta 74W 75Re 76Os 77Ir 78Pt 79Au 80Hg
7 89Ac 104Rf 105Db 106Sg 107Bh 108Hs 109Mt 110Ds 111Rg 112Cn

The elements of groups 4–11 are generally recognized as transition metals, justified by their typical chemistry, i.e. a large range of complex ions in various oxidation states, colored complexes, and catalytic properties either as the element or as ions (or both). Sc and Y in group 3 are also generally recognized as transition metals. However, the elements La–Lu and Ac–Lr and group 12 attract different definitions from different authors.

1. Many chemistry textbooks and printed periodic tables classify La and Ac as group 3 elements and transition metals, since their atomic ground-state configurations are s2d1 like Sc and Y. The elements Ce–Lu are considered as the "lanthanide" series (or "lanthanoid" according to IUPAC) and Th–Lr as the "actinide" series.[8]</ref>[9] The two series together are classified as f-block elements, or (in older sources) as "inner transition elements".
2. Some inorganic chemistry textbooks include La with the lanthanides and Ac with the actinides.[4][10][11] This classification is based on similarities in chemical behaviour and defines 15 elements in each of the two series, even though they correspond to the filling of an f subshell, which can only contain 14 electrons.
3. A third classification defines the f-block elements as La–Yb and Ac–No, while placing Lu and Lr in group 3.[5] This is based on the Aufbau principle (or Madelung rule) for filling electron subshells, in which 4f is filled before 5d (and 5f before 6d), so that the f subshell is actually full at Yb (and No), while Lu (and Lr) has an [ ]s2f14d1 configuration. However La and Ac are exceptions to the Aufbau principle with electron configuration [ ]s2d1 (not [ ]s2f1 as the Aufbau principle predicts), so it is not clear from atomic electron configurations whether La or Lu (Ac or Lr) should be considered as transition metals.[12]

Zinc, cadmium, and mercury are generally excluded from the transition metals,[5] as they have the electronic configuration [ ]d10s2, with no incomplete d shell.[13] In the oxidation state +2 the ions have the electronic configuration [ ]…d10. However, these elements can exist in other oxidation states, including the +1 oxidation state, as in the diatomic ion Hg2+
2
. The group 12 elements Zn, Cd and Hg may therefore, under certain criteria, be classed as post-transition metals in this case. However, it is often convenient to include these elements in a discussion of the transition elements. For example, when discussing the crystal field stabilization energy of first-row transition elements, it is convenient to also include the elements calcium and zinc, as both Ca2+
and Zn2+
have a value of zero, against which the value for other transition metal ions may be compared. Another example occurs in the Irving–Williams series of stability constants of complexes.

The recent (though disputed and so far not reproduced independently) synthesis of mercury(IV) fluoride (HgF
4
) has been taken by some to reinforce the view that the group 12 elements should be considered transition metals,[14] but some authors still consider this compound to be exceptional.[15]

Although meitnerium, darmstadtium, and roentgenium are within the d-block and are expected to behave as transition metals analogous to their lighter congeners iridium, platinum, and gold, this has not yet been experimentally confirmed.

Subclasses

Early transition metals are on the left side of the periodic table from group 3 to group 7. Late transition metals are on the right side of the d-block, from group 8 to 11 (and 12 if it is counted as transition metals).

Electronic configuration

The general electronic configuration of the d-block elements is [Inert gas] (n − 1)d1–10n s0–2. The period 6 and 7 transition metals also add (n − 2)f0–14 electrons, which are omitted from the tables below.

The Madelung rule predicts that the typical electronic structure of transition metal atoms can be written as [inert gas]ns2(n − 1)dm where the inner d orbital is predicted to be filled after the valence-shell's s orbital is filled. This rule is however only approximate – it only holds for some of the transition elements, and only then in their neutral ground state.

The d-sub-shell is the next-to-last sub-shell and is denoted as ${\displaystyle (n-1)d}$-sub-shell. The number of s electrons in the outermost s sub-shell is generally one or two except palladium (Pd), with no electron in that s-sub shell in its ground state. The s-sub-shell in the valence shell is represented as the ns sub-shell, e.g. 4s. In the periodic table, the transition metals are present in eight groups (4 to 11), with some authors including some elements in groups 3 or 12.

The elements in group 3 have an ns2(n − 1)d1 configuration. The first transition series is present in the 4th period, and starts after Ca (Z = 20) of group-2 with the configuration [Ar]4s2, or scandium (Sc), the first element of group 3 with atomic number Z = 21 and configuration [Ar]4s23d1, depending on the definition used. As we move from left to right, electrons are added to the same d-sub-shell till it is complete. The element of group 11 in the first transition series is copper (Cu) with an atypical configuration [Ar]4s13d10. Despite the filled d subshell in metallic copper it nevertheless forms a stable ion with an incomplete d subshell. Since the electrons added fill the ${\displaystyle (n-1)d}$ orbitals, the properties of the d-block elements are quite different from those of s and p block elements in which the filling occurs either in s or in p-orbitals of the valence shell. The electronic configuration of the individual elements present in all the d-block series are given below:[16]

 Group Atomic nr Element Electronconfiguration 3 4 5 6 7 8 9 10 11 12 21 22 23 24 25 26 27 28 29 30 Sc Ti V Cr Mn Fe Co Ni Cu Zn 3d14s2 3d24s2 3d34s2 3d54s1 3d54s2 3d64s2 3d74s2 3d84s2 3d104s1 3d104s2
 Atomic nr Element Electronconfiguration 39 40 41 42 43 44 45 46 47 48 Y Zr Nb Mo Tc Ru Rh Pd Ag Cd 4d15s2 4d25s2 4d45s1 4d55s1 4d55s2 4d75s1 4d85s1 4d10 4d105s1 4d105s2
 Atomic nr Element Electronconfiguration 57 72 73 74 75 76 77 78 79 80 La Hf Ta W Re Os Ir Pt Au Hg 5d16s2 5d26s2 5d36s2 5d46s2 5d56s2 5d66s2 5d76s2 5d96s1 5d106s1 5d106s2
 Atomic nr Element Electronconfiguration 89 104 105 106 107 108 109 110 111 112 Ac Rf Db Sg Bh Hs Mt Ds Rg Cn 6d17s2 6d27s2 6d37s2 6d47s2 6d57s2 6d67s2 6d77s2 6d87s2 6d97s2 6d107s2

A careful look at the electronic configuration of the elements reveals that there are certain exceptions, for example Cr and Cu. These are either because of the symmetry or nuclear-electron and electron-electron force.

The ${\displaystyle (n-1)d}$ orbitals that are involved in the transition metals are very significant because they influence such properties as magnetic character, variable oxidation states, formation of colored compounds etc. The valence ${\displaystyle s(ns)}$ and ${\displaystyle p(np)}$ orbitals have very little contribution in this regard since they hardly change in the moving from left to the right in a transition series. In transition metals, there is a greater horizontal similarities in the properties of the elements in a period in comparison to the periods in which the d-orbitals are not involved. This is because in a transition series, the valence shell electronic configuration of the elements do not change. However, there are some group similarities as well.

Characteristic properties

There are a number of properties shared by the transition elements that are not found in other elements, which results from the partially filled d shell. These include

• the formation of compounds whose color is due to dd electronic transitions
• the formation of compounds in many oxidation states, due to the relatively low energy gap between different possible oxidation states[17]
• the formation of many paramagnetic compounds due to the presence of unpaired d electrons. A few compounds of main group elements are also paramagnetic (e.g. nitric oxide, oxygen)

Most transition metals can be bound to a variety of ligands, allowing for a wide variety of transition metal complexes.[18]

Coloured compounds

From left to right, aqueous solutions of: Co(NO
3
)
2
(red); K
2
Cr
2
O
7
(orange); K
2
CrO
4
(yellow); NiCl
2
(turquoise); CuSO
4
(blue); KMnO
4
(purple).

Colour in transition-series metal compounds is generally due to electronic transitions of two principal types.

• charge transfer transitions. An electron may jump from a predominantly ligand orbital to a predominantly metal orbital, giving rise to a ligand-to-metal charge-transfer (LMCT) transition. These can most easily occur when the metal is in a high oxidation state. For example, the colour of chromate, dichromate and permanganate ions is due to LMCT transitions. Another example is that mercuric iodide, HgI2, is red because of a LMCT transition.

A metal-to-ligand charge transfer (MLCT) transition will be most likely when the metal is in a low oxidation state and the ligand is easily reduced.

In general charge transfer transitions result in more intense colours than d-d transitions.

• d-d transitions. An electron jumps from one d-orbital to another. In complexes of the transition metals the d orbitals do not all have the same energy. The pattern of splitting of the d orbitals can be calculated using crystal field theory. The extent of the splitting depends on the particular metal, its oxidation state and the nature of the ligands. The actual energy levels are shown on Tanabe-Sugano diagrams.

In centrosymmetric complexes, such as octahedral complexes, d-d transitions are forbidden by the Laporte rule and only occur because of vibronic coupling in which a molecular vibration occurs together with a d-d transition. Tetrahedral complexes have somewhat more intense colour because mixing d and p orbitals is possible when there is no centre of symmetry, so transitions are not pure d-d transitions. The molar absorptivity (ε) of bands caused by d-d transitions are relatively low, roughly in the range 5-500 M−1cm−1 (where M = mol dm−3).[19] Some d-d transitions are spin forbidden. An example occurs in octahedral, high-spin complexes of manganese(II), which has a d5 configuration in which all five electron has parallel spins; the colour of such complexes is much weaker than in complexes with spin-allowed transitions. Many compounds of manganese(II) appear almost colourless. The spectrum of [Mn(H
2
O)
6
]2+
shows a maximum molar absorptivity of about 0.04 M−1cm−1 in the visible spectrum.

Oxidation states

A characteristic of transition metals is that they exhibit two or more oxidation states, usually differing by one. For example, compounds of vanadium are known in all oxidation states between −1, such as [V(CO)
6
]
, and +5, such as VO3−
4
.

Main group elements in groups 13 to 18 also exhibit multiple oxidation states. The "common" oxidation states of these elements typically differ by two. For example, compounds of gallium in oxidation states +1 and +3 exist in which there is a single gallium atom. No compound of Ga(II) is known: any such compound would have an unpaired electron and would behave as a free radical and be destroyed rapidly. The only compounds in which gallium has a formal oxidation state of +2 are dimeric compounds, such as [Ga
2
Cl
6
]2−
, which contain a Ga-Ga bond formed from the unpaired electron on each Ga atom.[20] Thus the main difference in oxidation states, between transition elements and other elements is that oxidation states are known in which there is a single atom of the element and one or more unpaired electrons.

The maximum oxidation state in the first row transition metals is equal to the number of valence electrons from titanium (+4) up to manganese (+7), but decreases in the later elements. In the second row the maximum occurs with ruthenium (+8), and in the third row, the maximum occurs with iridium (+9). In compounds such as [MnO
4
]
and OsO
4
the elements achieve a stable configuration by covalent bonding.

The lowest oxidation states are exhibited in metal carbonyl complexes such as Cr(CO)
6
(oxidation state zero) and [Fe(CO)
4
]2−
(oxidation state −2) in which the 18-electron rule is obeyed. These complexes are also covalent.

Ionic compounds are mostly formed with oxidation states +2 and +3. In aqueous solution the ions are hydrated by (usually) six water molecules arranged octahedrally.

Magnetism

Transition metal compounds are paramagnetic when they have one or more unpaired d electrons.[21] In octahedral complexes with between four and seven d electrons both high spin and low spin states are possible. Tetrahedral transition metal complexes such as [FeCl
4
]2−
are high spin because the crystal field splitting is small so that the energy to be gained by virtue of the electrons being in lower energy orbitals is always less than the energy needed to pair up the spins. Some compounds are diamagnetic. These include octahedral, low-spin, d6 and square-planar d8 complexes. In these cases, crystal field splitting is such that all the electrons are paired up.

Ferromagnetism occurs when individual atoms are paramagnetic and the spin vectors are aligned parallel to each other in a crystalline material. Metallic iron and the alloy alnico are examples of ferromagnetic materials involving transition metals. Anti-ferromagnetism is another example of a magnetic property arising from a particular alignment of individual spins in the solid state.

Catalytic properties

The transition metals and their compounds are known for their homogeneous and heterogeneous catalytic activity. This activity is ascribed to their ability to adopt multiple oxidation states and to form complexes. Vanadium(V) oxide (in the contact process), finely divided iron (in the Haber process), and nickel (in catalytic hydrogenation) are some of the examples. Catalysts at a solid surface (nanomaterial-based catalysts) involve the formation of bonds between reactant molecules and atoms of the surface of the catalyst (first row transition metals utilize 3d and 4s electrons for bonding). This has the effect of increasing the concentration of the reactants at the catalyst surface and also weakening of the bonds in the reacting molecules (the activation energy is lowered). Also because the transition metal ions can change their oxidation states, they become more effective as catalysts.

An interesting type of catalysis occurs when the products of a reaction catalyse the reaction producing more catalyst (autocatalysis). One example is the reaction of oxalic acid with acidified potassium permanganate (or manganate (VII)).[22] Once a little Mn2+ has been produced, it can react with MnO4 forming Mn3+. This then reacts with C2O4 ions forming Mn2+ again.

Physical properties

As implied by the name, all transition metals are metals and thus conductors of electricity.

In general, transition metals possess a high density and high melting points and boiling points. These properties are due to metallic bonding by delocalized d electrons, leading to cohesion which increases with the number of shared electrons. However the group 12 metals have much lower melting and boiling points since their full d subshells prevent d–d bonding, which again tends to differentiate them from the accepted transition metals. Mercury has a melting point of −38.83 °C (−37.89 °F) and is a liquid at room temperature.

References

1. ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006–) "transition element". doi:10.1351/goldbook.T06456
2. ^ Petrucci, Ralph H.; Harwood, William S.; Herring, F. Geoffrey (2002). General chemistry: principles and modern applications (8th ed.). Upper Saddle River, N.J: Prentice Hall. pp. 341–342. ISBN 978-0-13-014329-7. LCCN 2001032331. OCLC 46872308.
3. ^ Housecroft, C. E. and Sharpe, A. G. (2005) Inorganic Chemistry, 2nd ed, Pearson Prentice-Hall, pp. 20–21.
4. ^ a b Cotton, F. A. and Wilkinson, G. (1988) Inorganic Chemistry, 5th ed., Wiley, pp. 625–627. ISBN 978-0-471-84997-1.
5. ^ a b c Jensen, William B. (2003). "The Place of Zinc, Cadmium, and Mercury in the Periodic Table" (PDF). Journal of Chemical Education. 80 (8): 952–961. Bibcode:2003JChEd..80..952J. doi:10.1021/ed080p952.
6. ^ Bury, C. R. (1921). "Langmuir's theory of the arrangement of electrons in atoms and molecules". J. Am. Chem. Soc. 43 (7): 1602–1609. doi:10.1021/ja01440a023.
7. ^ Bury, Charles Rugeley. Encyclopedia.com Complete dictionary of scientific biography (2008).
8. ^ Petrucci, Harwood & Herring 2002, pp. 49–50, 951.
9. ^ Miessler, G. L. and Tarr, D. A. (1999) Inorganic Chemistry, 2nd edn, Prentice-Hall, p. 16. ISBN 978-0-13-841891-5.
10. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
11. ^ Housecroft, C. E. and Sharpe, A. G. (2005) Inorganic Chemistry, 2nd ed., Pearson Prentice-Hall, p. 741.
12. ^ Scerri, E. R. (2011) A Very Short Introduction to the Periodic Table, Oxford University Press.
13. ^ Cotton, F. Albert; Wilkinson, G.; Murillo, C. A. (1999). Advanced Inorganic Chemistry (6th ed.). New York: Wiley, ISBN 978-0-471-19957-1.
14. ^ Wang, Xuefang; Andrews, Lester; Riedel, Sebastian; Kaupp, Martin (2007). "Mercury Is a Transition Metal: The First Experimental Evidence for HgF4". Angew. Chem. Int. Ed. 46 (44): 8371–8375. doi:10.1002/anie.200703710. PMID 17899620.
15. ^ Jensen, William B. (2008). "Is Mercury Now a Transition Element?". J. Chem. Educ. 85 (9): 1182–1183. Bibcode:2008JChEd..85.1182J. doi:10.1021/ed085p1182.
16. ^ Miessler, G. L. and Tarr, D. A. (1999) Inorganic Chemistry, 2nd edn, Prentice-Hall, p. 38 ISBN 978-0-13-841891-5
17. ^ Matsumoto, Paul S (2005). "Trends in Ionization Energy of Transition-Metal Elements". Journal of Chemical Education. 82 (11): 1660. Bibcode:2005JChEd..82.1660M. doi:10.1021/ed082p1660.
18. ^ Hogan, C. Michael (2010). "Heavy metal" in Encyclopedia of Earth. National Council for Science and the Environment. E. Monosson and C. Cleveland (eds.) Washington DC.
19. ^ Orgel, L.E. (1966). An Introduction to Transition-Metal Chemistry, Ligand field theory (2nd. ed.). London: Methuen.
20. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8. p. 240
21. ^ Figgis, B.N.; Lewis, J. (1960). Lewis, J.; Wilkins, R.G. (eds.). The Magnetochemistry of Complex Compounds. Modern Coordination Chemistry. New York: Wiley Interscience. pp. 400–454.
22. ^ Kovacs KA, Grof P, Burai L, Riedel M (2004). "Revising the Mechanism of the Permanganate/Oxalate Reaction". J. Phys. Chem. A. 108 (50): 11026–11031. Bibcode:2004JPCA..10811026K. doi:10.1021/jp047061u.
Block (periodic table)

A block of the periodic table is a set of chemical elements predominantly characterised by having their highest energy electrons in the same atomic orbital type. The term appears to have been first used by Charles Janet. Each block is named after its characteristic orbital; thus, the blocks are:

s-block

p-block

d-block

f-blockThe block names (s, p, d, f and g) are derived from the spectroscopic notation for the associated atomic orbitals: sharp, principal, diffuse, and fundamental.

Carbide

In chemistry, a carbide is a compound composed of carbon and a less electronegative element. Carbides can be generally classified by the chemical bonds type as follows: (i) salt-like, (ii) covalent compounds, (iii) interstitial compounds, and (iv) "intermediate" transition metal carbides. Examples include calcium carbide (CaC2), silicon carbide (SiC), tungsten carbide (WC; often called, simply, carbide when referring to machine tooling), and cementite (Fe3C), each used in key industrial applications. The naming of ionic carbides is not systematic.

Chalcogenide

A chalcogenide is a chemical compound consisting of at least one chalcogen anion and at least one more electropositive element. Although all group 16 elements of the periodic table are defined as chalcogens, the term chalcogenide is more commonly reserved for sulfides, selenides, tellurides, and polonides, rather than oxides. Many metal ores exist as chalcogenides. Photoconductive chalcogenide glasses are used in xerography. Some pigments and catalysts are also based on chalcogenides. The metal dichalcogenide MoS2 is a common solid lubricant.

Cluster chemistry

In chemistry, a cluster is an ensemble of bound atoms or molecules that is intermediate in size between a molecule and a bulk solid. Clusters exist of diverse stoichiometries and nuclearities. For example, carbon and boron atoms form fullerene and borane clusters, respectively. Transition metals and main group elements form especially robust clusters. Clusters can also consist solely of a certain kind of molecules, such as water clusters.

The phrase cluster was coined by F.A. Cotton in the early 1960s to refer to compounds containing metal–metal bonds. In another definition a cluster compound contains a group of two or more metal atoms where direct and substantial metal bonding is present. The prefixed terms "nuclear" and "metallic" are used and imply different meanings. For example, polynuclear refers to a cluster with more than one metal atom, regardless of the elemental identities. Heteronuclear refers to a cluster with at least two different metal elements.

The main cluster types are "naked" clusters (without stabilizing ligands) and those with ligands. For transition metal clusters, typical stabilizing ligands include carbon monoxide, halides, isocyanides, alkenes, and hydrides. For main group elements, typical clusters are stabilized by hydride ligands.

Transition metal clusters are frequently composed of refractory metal atoms. In general metal centers with extended d-orbitals form stable clusters because of favorable overlap of valence orbitals. Thus, metals with a low oxidation state for the later metals and mid-oxidation states for the early metals tend to form stable clusters. Polynuclear metal carbonyls are generally found in late transition metals with low formal oxidation states. The polyhedral skeletal electron pair theory or Wade's electron counting rules predict trends in the stability and structures of many metal clusters. Jemmis mno rules have provided additional insight into the relative stability of metal clusters.

Coordination complex

In chemistry, a coordination complex consists of a central atom or ion, which is usually metallic and is called the coordination centre, and a surrounding array of bound molecules or ions, that are in turn known as ligands or complexing agents. Many metal-containing compounds, especially those of transition metals, are coordination complexes. A coordination complex whose centre is a metal atom is called a metal complex.

Group 7 element

Group 7, numbered by IUPAC nomenclature, is a group of elements in the periodic table. They are manganese (Mn), technetium (Tc), rhenium (Re), and bohrium (Bh). All known elements of group 7 are transition metals.

Like other groups, the members of this family show patterns in their electron configurations, especially the outermost shells resulting in trends in chemical behavior.

Group 8 element

Group 8 is a group (column) of chemical elements in the periodic table. It consists of iron (Fe), ruthenium (Ru), osmium (Os) and hassium (Hs). They are all transition metals.

Like other groups, the members of this family show patterns in electron configuration, especially in the outermost shells, resulting in trends in chemical behavior.

"Group 8" is the modern standard designation for this group, adopted by the IUPAC in 1990.In the older group naming systems, this group was combined with group 9 (cobalt, rhodium, iridium, and meitnerium) and group 10 (nickel, palladium, platinum, and darmstadtium) and called group "VIIIB" in the Chemical Abstracts Service (CAS) "U.S. system", or "VIII" in the old IUPAC (pre-1990) "European system" (and in Mendeleev's original table).

Group 8 (current IUPAC) should not be confused with "group VIIIA" in the CAS system, which is group 18 (current IUPAC), the noble gases.

While groups (columns) of the periodic table are sometimes named after their lighter member (as in "the oxygen group" for group 16), the term iron group does not mean "group 8". Most often, it means a set of adjacent elements on period (row) 4 of the table that includes iron, such as chromium, manganese, iron, cobalt, and nickel; or only the last three; or some other set — depending on the context.

Group 9 element

Group 9 is a group (column) of chemical elements in the periodic table. Members are cobalt (Co), rhodium (Rh), iridium (Ir) and meitnerium (Mt). These are all transition metals in the d-block.

Like other groups, the members of this family show patterns in electron configuration, especially in the outermost shells, resulting in trends in chemical behavior; however, rhodium deviates from the pattern.

"Group 9" is the modern standard designation for this group, adopted by the IUPAC in 1990.In the older group naming systems, this group was combined with group 8 (iron, ruthenium, osmium, and hassium) and group 10 (nickel, palladium, platinum, and darmstadtium) and called group "VIIIB" in the Chemical Abstracts Service (CAS) "U.S. system", or "VIII" in the old IUPAC (pre-1990) "European system" (and in Mendeleev's original table).

Inorganic chemistry

Inorganic chemistry deals with the synthesis and behavior of inorganic and organometallic compounds. This field covers all chemical compounds except the myriad organic compounds (carbon-based compounds, usually containing C-H bonds), which are the subjects of organic chemistry. The distinction between the two disciplines is far from absolute, as there is much overlap in the subdiscipline of organometallic chemistry. It has applications in every aspect of the chemical industry, including catalysis, materials science, pigments, surfactants, coatings, medications, fuels, and agriculture.

Osmium dioxide

Osmium dioxide is an inorganic compound with the formula OsO2. It exists as brown to black crystalline powder, but single crystals are golden and exhibit metallic conductivity. The compound crystallizes in the rutile structural motif, i.e. the connectivity is very similar to that in the mineral rutile.

Oxide

An oxide is a chemical compound that contains at least one oxygen atom and one other element in its chemical formula. "Oxide" itself is the dianion of oxygen, an O2– atom. Metal oxides thus typically contain an anion of oxygen in the oxidation state of −2. Most of the Earth's crust consists of solid oxides, the result of elements being oxidized by the oxygen in air or in water. Hydrocarbon combustion affords the two principal carbon oxides: carbon monoxide and carbon dioxide. Even materials considered pure elements often develop an oxide coating. For example, aluminium foil develops a thin skin of Al2O3 (called a passivation layer) that protects the foil from further corrosion. Individual elements can often form multiple oxides, each containing different amounts of the element and oxygen. In some cases these are distinguished by specifying the number of atoms as in carbon monoxide and carbon dioxide, and in other cases by specifying the element's oxidation number, as in iron(II) oxide and iron(III) oxide. Certain elements can form many different oxides, such as those of nitrogen.

Period (periodic table)

A period in the periodic table is a row of chemical elements. All elements in a row have the same number of electron shells. Each next element in a period has one more proton and is less metallic than its predecessor. Arranged this way, groups of elements in the same column have similar chemical and physical properties, reflecting the periodic law. For example, the alkali metals lie in the first column (group 1) and share similar properties, such as high reactivity and the tendency to lose one electron to arrive at a noble-gas electronic configuration. As of 2016, a total of 118 elements have been discovered and confirmed.

Modern quantum mechanics explains these periodic trends in properties in terms of electron shells. As atomic number increases, shells fill with electrons in approximately the order shown at right. The filling of each shell corresponds to a row in the table.

In the s-block and p-block of the periodic table, elements within the same period generally do not exhibit trends and similarities in properties (vertical trends down groups are more significant). However, in the d-block, trends across periods become significant, and in the f-block elements show a high degree of similarity across periods.

Period 4 element

A period 4 element is one of the chemical elements in the fourth row (or period) of the periodic table of the elements. The periodic table is laid out in rows to illustrate recurring (periodic) trends in the chemical behaviour of the elements as their atomic number increases: a new row is begun when chemical behaviour begins to repeat, meaning that elements with similar behaviour fall into the same vertical columns. The fourth period contains 18 elements, beginning with potassium and ending with krypton. As a rule, period 4 elements fill their 4s shells first, then their 3d and 4p shells, in that order; however, there are exceptions, such as chromium.

Period 5 element

A period 5 element is one of the chemical elements in the fifth row (or period) of the periodic table of the elements. The periodic table is laid out in rows to illustrate recurring (periodic) trends in the chemical behaviour of the elements as their atomic number increases: a new row is begun when chemical behaviour begins to repeat, meaning that elements with similar behaviour fall into the same vertical columns. The fifth period contains 18 elements, beginning with rubidium and ending with xenon. As a rule, period 5 elements fill their 5s shells first, then their 4d, and 5p shells, in that order; however, there are exceptions, such as rhodium.

Period 7 element

A period 7 element is one of the chemical elements in the seventh row (or period) of the periodic table of the chemical elements. The periodic table is laid out in rows to illustrate recurring (periodic) trends in the chemical behaviour of the elements as their atomic number increases: a new row is begun when chemical behaviour begins to repeat, meaning that elements with similar behaviour fall into the same vertical columns. The seventh period contains 32 elements, tied for the most with period 6, beginning with francium and ending with oganesson, the heaviest element currently discovered. As a rule, period 7 elements fill their 7s shells first, then their 5f, 6d, and 7p shells, in that order; however, there are exceptions, such as plutonium.

Post-transition metal

Post-transition metals are a set of metallic elements in the periodic table located between the transition metals to their left, and the metalloids to their right. Depending on where these adjacent groups are judged to begin and end, there are at least five competing proposals for which elements to include: the three most common contain six, ten and thirteen elements, respectively (see image). All proposals include gallium, indium, tin, thallium, lead, and bismuth.

Physically, post-transition metals are soft (or brittle), have poor mechanical strength, and have melting points lower than those of the transition metals. Being close to the metal-nonmetal border, their crystalline structures tend to show covalent or directional bonding effects, having generally greater complexity or fewer nearest neighbours than other metallic elements.

Chemically, they are characterised—to varying degrees—by covalent bonding tendencies, acid-base amphoterism and the formation of anionic species such as aluminates, stannates, and bismuthates (in the case of aluminium, tin, and bismuth, respectively). They can also form Zintl phases (half-metallic compounds formed between highly electropositive metals and moderately electronegative metals or metalloids).

The name is universally used, but not officially sanctioned by any organization such as the IUPAC. The origin of the term is unclear: one early use was in 1940 in a chemistry text. Alternate names for this group are B-subgroup metals, other metals, and p-block metals; and at least eleven other labels.

Rhenium disulfide

Rhenium disulfide is an inorganic compound of rhenium and sulfur with the formula ReS2. It has a layered structure where atoms are strongly bonded within each layer. The layers are held together by weak Van der Waals bonds, and can be easily peeled off from the bulk material. It is a two-dimensional (2D) group VII transition metal dichalcogenide (TMD). Different to other TMDs, ReS2 has shown layer-independent electrical, optical, and vibrational properties.Nanostructured ReS2 can usually be achieved through mechanical exfoliation, chemical vapor deposition (CVD), and chemical and liquid exfoliations. It is widely used in electronic and optoelectronic device, energy storage, photocatalytic and electrocatalytic reactions.

Transition metal carbene complex

A transition metal carbene complex is an organometallic compound featuring a divalent organic ligand. The divalent organic ligand coordinated to the metal center is called a carbene. Carbene complexes for almost all transition metals have been reported. Many methods for synthesizing them and reactions utilizing them have been reported. The term carbene ligand is a formalism since many are not derived from carbenes and almost none exhibit the reactivity characteristic of carbenes. Described often as M=CR2, they represent a class of organic ligands intermediate between alkyls (−CR3) and carbynes (≡CR). They feature in some catalytic reactions, especially alkene metathesis, and are of value in the preparation of some fine chemicals.

Tungsten trioxide

Tungsten(VI) oxide, also known as tungsten trioxide or tungstic anhydride, WO3, is a chemical compound containing oxygen and the transition metal tungsten. It is obtained as an intermediate in the recovery of tungsten from its minerals. Tungsten ores are treated with alkalis to produce WO3. Further reaction with carbon or hydrogen gas reduces tungsten trioxide to the pure metal.　Tungsten trioxide is a strong oxidative agent, it reacts rare-earth elements, iron, copper, aluminium, manganese, zinc, chromium, molybdenum, carbon, hydrogen and silver to make the pure tungsten metal, and gold and platinum to make the tungsten dioxide.

2 WO3 + 3 C → 2 W + 3 CO2 (high temperature)

WO3 + 3 H2 → W + 3 H2O (550 - 850 °C)

WO3 + 2Fe → W + Fe2O3

2WO3 + Pt → 2WO2 + PtO2Tungsten(VI) oxide occurs naturally in the form of hydrates, which include minerals: tungstite WO3·H2O, meymacite WO3·2H2O and hydrotungstite (of the same composition as meymacite, however sometimes written as H2WO4). These minerals are rare to very rare secondary tungsten minerals.

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