Is SCN one-toed or two-toed
The Complex chemistry(Coordination chemistry) is the field of inorganic chemistry that deals with chemical complex compounds.
A complex (or Coordination connection) is a structure in which a Central atom (usually a metal ion), which has gaps in its electron configuration, is surrounded by one or more molecules or ions (the ligands), each of which has at least one free electron pair provide for binding. This type of bond differs from the other forms of chemical bond (covalent bond, ionic bond, metal bond). One speaks of one Complex binding, a coordinative bond or, if the central atom is a metal, of one Metal complex.
The ligands enclose the central atom - the word complex derives from the Latin verb complecti or its participle complexus from, its German equivalent embrace, embrace is.
Complex compounds often play a decisive role in biology. The compounds hemoglobin and chlorophyll, which are extremely important for life, contain metal complexes.
Many complex compounds are colored and can therefore be used as dyes. Complex compounds are often accessible from the corresponding salts of a central ion. For example, anhydrous, white copper sulphate turns light blue when water (aqua) is added. An aqua complex of copper is created in which four water molecules act as ligands of the central ion (complex formation reaction):
The pioneer of coordination chemistry, Alfred Werner, received the Nobel Prize in Chemistry in 1913 for his correct interpretation of the structure and bonding relationships in complexes.
The complex formation reaction is a classic acid-base reaction according to the theory of Gilbert Newton Lewis. Here is the Central atom (mostly a metal cation, most often transition metals) is the Lewis acid (electron pair acceptor); the Lewis base is the Ligand, a particle or molecule that contains at least one free electron pair (Electron pair donor) can provide for bond formation. This type of chemical bond is called coordinative bond designated. Since several (at least two) ligands bind to a central atom in complexes, one also speaks of higher-order compounds.
The Central particles are often cations, but they can also be neutral or (rarely) anions:
The Ligands can be inorganic or organic in nature:
- inorganic ligands:
- organic ligands:
The nomenclature of complex compounds
Rules for the nomenclature: For the systematic naming of complex salts, the cation (regardless of whether it is complex or not) is given first and then the anion. The components of a coordination unit are named in the following order:
- The Number of ligands is indicated by preceding Greek numerals: mono, di, tri, tetra, penta, hexa, hepta, octa etc. For ligands with complicated names or to avoid ambiguities (e.g. dithiosulphate), the multipliers derived from the Greek are used: bis, tris, tetrakis, pentakis, hexakis, heptakis, octakis etc. The part multiplied by this comes in brackets.
- Type of ligands: The various ligands are listed in alphabetical order, regardless of their number or charge. Anionic ligands are given the ending -o on their anion names (e.g. chlorido). The names of neutral or cationic ligands are not changed. Exceptions to this rule are the names of water (aqua), ammonia (ammin), CO (carbonyl) and NO (nitrosyl).
- Central ion: In a complex anion, the central ion (with the Latin root word) has the ending -at. If the complex is a cation or a neutral molecule, then the name of the central ion (with German name) does not change,
- Charge of the central ion: The charge of the central ion (= oxidation number) is indicated by a Roman number ("stick number") in round brackets and placed after the name of the coordination unit. (A plus sign is not written; the Arabic number 0 is used for zero.)
The full name of the coordination unit is written in one word. Except for the names of the ligands aqua, ammin and nitrosyl, the names of all neutral ligands are put in brackets. The names of inorganic anionic ligands are put in round brackets if they already contain numerical prefixes or if this avoids ambiguity. In the name of complex salts, a hyphen is written between the names of the cation and the anion.
- K3[Fe (CN)6] Potassium hexacyanoferrate (III).
- Write coordination unit in square brackets, if charge is present, as exponent
- Central atom in front of ligands
- anionic ligands before neutral ligands
- Polyatomic ligands in parentheses
Name of complex compounds
- Ligands in alphabetical order before the name of the central atom. Multiple ligands are given the following (Greek) prefixes: di (2), tri (3), tetra (4), penta (5), hexa (6), hepta (7), octa (8), nona (9).
- Anionic ligands are given the ending "-o".
- Important anionic ligands: '
- F.- (fluorido); Cl- (chlorido); Br- (bromido); I.- (iodido)
- The name is the ion name + o
- O2- (oxido); O22- (peroxido); OH- (hydroxido); H- (hydrido)
- S.2- (thio, sulfido); SO42- (sulfato); S.2O32- (thiosulfato); NO3- (nitrato)
- NO2- (nitrito, nitro with coordination via N or nitrito-N and nitrito-O)
- CN- (cyanido, isocyanido with coordination via N or cyanido-C and cyanido-N)
- SCN- (thiocyanato or isothiocyanato when coordinated via N)
- Important neutral ligands:
- NH3 (ammin); H2O (aqua, obsolete aquo); CO (carbonyl); NO (nitrosyl)
- If the entire coordination unit is an anion, it is given the ending -at. The Latin name is used for the central atom (e.g. argentate, ferrate, cuprate, etc.).
- If the coordination unit is neutral or a cation, the unchanged German name is used.
- The oxidation number (Roman numeral) of the central atom comes after the name of the central atom.
- [Fe (CN)6]3- Hexacyanidoferrate (III);
- [Cu (NH3)4]2+ Tetraammine copper (II);
- [CrCl3(H2O)3] Triaquatrichloridochrome (III);
- [FeBr2(CN)2(H2O)2]2- Diaquadibromidodicyanidoferrate (II)
Determination of the oxidation number, valence electron number and estimation of the stability
The oxidation number of the central particle is determined by considering the total charge on the complex and the charges on the ligands. The sum of the charge contributions of the ligands and the oxidation number of the central particle (s) must give the charge of the complex.
- Single negatively charged ligands: e.g. B. Cl-, Br-, Alkyl, hydride, Cp;
- Neutral ligands: e.g. B. (Ph)3P, CO, C6H6, Butadiene
The number of valence electrons is the sum of the electrons of the ligands and the central particle to which the ligands are coordinated.
- Central particles: e.g. B. Fe- (0) brings 8 electrons into the complex because it is in the 8th subgroup, Fe- (II) therefore has 6 in its d orbitals.
- Ligands: e.g. B. bring Cl- and (Ph)3P two electrons with one, η5-Cp and η6-C6H6 six electrons, unbridged μ1-CO 2 electrons, bridged μ2-CO one electron (ηn, μn: see hapticity)
An estimate of the stability can be made with the 18-electron rule, if this fails, one must use the ligand field theory or the molecular orbital theory.
Geometry of Complexes
The Coordination number indicates with how many so-called monodentate ligands a central atom surrounds itself. Free electron pairs are not to be neglected here. The coordination numbers are particularly common 2, 4 and 6.
If the coordination number is the same:
- two, there is a linear complex
- three either a trigonal-planar or a trigonal-aplanar structure is obtained (the central particle is not exactly in the middle of the triangle, but slightly above it)
- four the ligands result in a tetrahedron or a square-planar structure
- five the result is a square-pyramidal or trigonal-bipyramidal structure. Both can be converted into one another by the Berry pseudo-rotation and are in equilibrium at the appropriate temperature.
- six the ligands result in an octahedron or a trigonal antiprism or (less often) a trigonal prism
- seven (very rarely) you get a pentagonal bipyramid or a simply capped octahedron
- eight the ligands form a cube, a square antiprism or a trigonododecahedron
Only the coordination number is worth mentioning 12that gives an icosahedron or a cuboctahedron.
Symmetry of complexes
see group theory
Color of complexes
Complex compounds are often colored because they have larger delocalized electron systems. The charge transfer complexes show particularly intense colors, such as B. the permanganate. See also ligand field theory
Chelate complexes and toughness
The dentition indicates how many bonds to the central atom a ligand can form. Ligands that only bond to the central atom are called monodentate or monodentate. Ammonia (NH3, in the complex as Ammin for example, a monodentate ligand: H3N-M.
If a ligand has several coordination points that can be used simultaneously for coordination at the same metal center, one speaks of a chelate ligand (Greek chelé = Crab claw). These chelate complexes have a higher stability both thermodynamically and kinetically. The high thermodynamic stability is based on the increase in the entropy of the system, since the following reaction takes place in aqueous solution to form an octahedral complex, for example, with a bidentate ligand (ligand with two coordination sites):
Here four free particles (on the left) become seven free particles (on the right). The kinetic stability is based on the fact that for the formation of the complex (according to the kinetic gas theory) Less Particles have to hit and all bonds of a ligand to the central atom have to be opened simultaneously during the dissociation.
Examples of chelating ligands:
- An example of a bidentate ligand is ethylenediamine (C.2H8N2) (Structure see figure)
- tetradentate ligands: e.g. B. NTA: nitrilotriacetic acid
- hexadentate ligands: e.g. B. EDTA: ethylene diamine tetraacetate, (−OOC-H2C–)2N-CH2–CH2–N (–CH2-COO−)2. EDTA can be used to soften water because it reacts with calcium to form easily soluble chelate complexes.
- Other important multidentate ligands are bipyridine and phenanthroline.
Polynuclear complexes contain more than one central atom. They are linked by a bridging ligand, for example oxygen or chlorine. Often it is a multi-electron center bond.Cl \ / \ / RhRh / \ / \ Cl
However, there are also complex compounds with (partly non-integer) metal-metal multiple bonds, e.g. B. [Tc2X9]3-, X = Cl, Br.
Hard and Soft Lewis Acids and Bases (HSAB)
The concept of hard and soft Lewis acids and bases (Hard and S.often A.cids and B.ases) was introduced by Pearson in 1963.
The hardness of an acid increases with decreasing size, increasing charge and decreasing polarizability of the acid particles. Bases are harder, the smaller, less polarizable and more difficult to oxidize the base particles.
Examples of Lewis acids:
- Hart: Fe3+, Al3+, Approx2+, Ti4+
- Transition area: Fe2+, Cu2+, Pb2+, Zn2+
- Soft: Au+, Cu+, Cd2+, Tl+
Examples of Lewis bases:
- Hart: F−, OH−, O2−, H2O, NH3
- Transitional area: Br−, N3−, NO2−
- Soft: I.−, S.2−, SCN−
Reactions of hard acids with hard bases and of soft acids with soft bases lead to more stable connections than the combinations soft - hard.
Application of the law of mass action
The Lewis acid-base reactions for complex formation are equilibrium reactions for which the law of mass action can be established. The overall reaction can be divided into individual steps (so-called elementary reactions), i. H. each for the addition of a ligand. The product of the equilibrium constants of the individual elementary reactions for complex formation then gives the equilibrium constant for the overall reaction.
The resulting constant is called Complex formation constant. This constant also indicates how stable the complex is or whether it tends to dissociate. Therefore, the complex formation constant also becomes Complex stability constant or Complex association constant KA. called. Their reciprocal value is called Complex dissociation constant KD. referred to, i.e. KA.-1 = KD.. The higher the complex formation constant, the more stable the complex, the smaller, the easier the dissociation.
The bond between the central atom and the ligand can be explained more or less comprehensively by different models
- Valence structure theory (valence bond theory, VB theory): ligand orbitals overlap with unoccupied hybrid orbitals of the central atom. The VB theory explains the geometry quite well, but z. B. not the colourfulness of complexes.
- Crystal field theory: The crystal field theory is based on pure electrostatic interactions between the ligands and the central atom. It explains the colourfulness of the complexes.
- Ligand field theory: The ligand field theory is an extension of the crystal field theory. She investigates the influence of the point-like ligands on the energies of the d orbitals of the central metal. (See also: Jahn-Teller effect).
- Molecular orbital theory: The molecular orbital theory provides the best description of complex compounds as it treats both the central atom and the ligands in a quantum-mechanical way.
Application and meaning
Complexes also play an important role in biology. These can be catalytically active proteins (enzymes) or catalytically inactive proteins. Many enzymes contain complexes in their active centers. This topic is one of the main areas of bioinorganic chemistry. In general, a complexing metal atom is present here, which is not completely complexed by amino acid side chains as ligands. A ligand site acts as an active center for the conversion or temporary binding of the substrate. The most common complex centers are iron, copper, zinc, calcium, magnesium and manganese. But there are also more unusual elements such as vanadium. Calcium in particular, as well as zinc complexes, are of structural importance (e.g. zinc fingers in DNA sequence recognition). The catalytically inactive proteins include z. B. Porphyrin complexes such as the heme in hemoglobin and in cytochromes, or chlorophyll (chelate complexes in each case). See also:
Various Complexing agents serve as food additives:
- Iminodisuccinate tetrasodium salt (This particularly biodegradable Complexing agents is used in water cycles to prevent and dissolve limescale deposits.)
In analytical chemistry, complex formation reactions with certain complexing agents are important as detection reactions (for copper, silver, nitrate / ring samples, bismuth ions). See also chelatometry.
Phthalocyanine complexes are used in CDs as a storage medium.
- Henry Taube: Electron Transfer Between Metal Complexes - A Review (Nobel Lecture). Angewandte Chemie 96 (5), pp. 315-326 (1984), ISSN 0044-8249
Categories: Subfield of Chemistry | Chemical bond | Food chemistry
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