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GREENWOOD / BLACK WALL STREET WILMINGTON 1898 LITTLE ROCK NINE TUSKEGEE STUDY ROSEWOOD 1923 THE TUSKEGEE AIRMEN SENECA VILLAGE BLACK TOWNS GREENWOOD / BLACK WALL STREET WILMINGTON 1898 LITTLE ROCK NINE TUSKEGEE STUDY ROSEWOOD 1923 THE TUSKEGEE AIRMEN SENECA VILLAGE BLACK TOWNS
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MY'STORY

The MOVE Fire

This is a personal recollection on the Move fire on May 13, 1985 Philadelphia police fired thousands of rounds at the MOVE house, city officials approved dropping an explosive device on the roof, the resulting fire was allowed to burn, 11 people—including five children—died, and 61 homes were destroyed. Philadelphia City Council later called it a “brutal attack carried out by the City of Philadelphia on its own citizens” and acknowledged that no individual faced criminal consequences for the bombing. One timeline correction worth preserving for the BHP record: the major previous MOVE-police confrontation was August 8, 1978, about seven years before the bombing, not a year or two earlier. Officer James Ramp was killed, other police and firefighters were wounded, nine MOVE members were later convicted, and television cameras recorded police beating Delbert Africa during his arrest. The 1985 MOVE Commission later specifically criticized city planners for failing to adequately use lessons from that 1978 confrontation. And that actually strengthens the point you’re making: 1985 did not happen without precedent or institutional memory. There had already been a deadly confrontation with MOVE, years of conflict, negotiations and police involvement before Osage Avenue.

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BLACK FACTS
The Truths They Never Taught You...

Shirley Chisholm — Unbought and Unbossed

In 1968 Shirley Chisholm became the first Black woman elected to the United States Congress. In 1972 she launched a campaign for the Democratic presidential nomination, breaking another political barrier.

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BHP gathered finds from its connected research sources. Showing the 4 strongest Black History matches.
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Wikipedia

Pi backbonding

In chemistry, pi backbonding or π backbonding is a π-bonding interaction between a filled (or half filled) orbital of a transition metal atom and a vacant orbital on an adjacent ion or molecule.[1][2] In this type of interaction, electrons from the metal are used to bond to the ligand, which dissipates excess negative charge and stabilizes the metal. It is common in transition metals with low oxidation states that have ligands such as carbon monoxide, olefins, or phosphines. The ligands involved in π backbonding can be broken into three groups: carbonyls and nitrogen analogs, alkenes and alkynes, and phosphines. Compounds where π backbonding is prominent include Ni(CO)4, Zeise's salt, and molybdenum and iron dinitrogen complexes.

Metal carbonyls, nitrosyls, and isocyanides

[edit]
σ bonding from electrons in CO's HOMO to metal center d-orbital.
π backbonding from electrons in metal center d-orbital to CO's LUMO.

The electrons are partially transferred from a d-orbital of the metal to anti-bonding molecular orbitals of CO (and its analogs). This electron-transfer strengthens the metal–C bond and weakens the C–O bond. The strengthening of the M–CO bond is reflected in increases of the vibrational frequencies for the M–C bond (often outside of the range for the usual IR spectrophotometers). Furthermore, the M–CO bond length is shortened. The weakening of the C–O bond is indicated by a decrease in the wavenumber of the νCO band(s) from that for free CO (2143 cm−1), for example to 2060 cm−1 in Ni(CO)4 and 1981 cm−1 in Cr(CO)6, and 1790 cm−1 in the anion [Fe(CO)4]2−.[3] For this reason, IR spectroscopy is an important diagnostic technique in metal–carbonyl chemistry. The article infrared spectroscopy of metal carbonyls discusses this in detail.

Many ligands other than CO are strong "backbonders". Nitric oxide is an even stronger π-acceptor than CO and νNO is a diagnostic tool in metal–nitrosyl chemistry. Isocyanides, RNC, are another class of ligands that are capable of π-backbonding. In contrast with CO, the σ-donor lone pair on the C atom of isocyanides is antibonding in nature and upon complexation the CN bond is strengthened and the νCN increased. At the same time, π-backbonding lowers the νCN. Depending on the balance of σ-bonding versus π-backbonding, the νCN can either be raised (for example, upon complexation with weak π-donor metals, such as Pt(II)) or lowered (for example, upon complexation with strong π-donor metals, such as Ni(0)).[4] For the isocyanides, an additional parameter is the MC=N–C angle, which deviates from 180° in highly electron-rich systems. Other ligands have weak π-backbonding abilities, which creates a labilization effect of CO, which is described by the cis effect.

Metal–alkene and metal–alkyne complexes

[edit]
σ bonding from electrons in alkene's HOMO to metal center d-orbital.
π backbonding from electrons in metal center d-orbital to alkene's LUMO.

As in metal–carbonyls, electrons are partially transferred from a d-orbital of the metal to antibonding molecular orbitals of the alkenes and alkynes.[5][6] This electron transfer strengthens the metal–ligand bond and weakens the C–C bonds within the ligand.[7] In the case of metal-alkenes and alkynes, the strengthening of the M–C2R4 and M–C2R2 bond is reflected in bending of the C–C–R angles which assume greater sp3 and sp2 character, respectively.[8][6] Thus strong π backbonding causes a metal-alkene complex to assume the character of a metallacyclopropane.[5] Alkenes and alkynes with electronegative substituents exhibit greater π backbonding.[6] Some strong π backbonding ligands are tetrafluoroethylene, tetracyanoethylene, and hexafluoro-2-butyne. In case of d10 configured complexes without vacant valence d-orbitals, strong π backbonding renders olefin and arene ligands effectively Z-type acceptor ligands.[9]

Metal-phosphine complexes

[edit]
R3P–M σ bonding
R3P–M π backbonding

Phosphines accept electron density from metal p or d orbitals into combinations of P–C σ* antibonding orbitals that have π symmetry.[10] When phosphines bond to electron-rich metal atoms, backbonding would be expected to lengthen P–C bonds as P–C σ* orbitals become populated by electrons. The expected lengthening of the P–C distance is often hidden by an opposing effect: as the phosphorus lone pair is donated to the metal, P(lone pair)–R(bonding pair) repulsions decrease, which acts to shorten the P–C bond. The two effects have been deconvoluted by comparing the structures of pairs of metal-phosphine complexes that differ only by one electron.[11] Oxidation of R3P–M complexes results in longer M–P bonds and shorter P–C bonds, consistent with π-backbonding.[12] In early work, phosphine ligands were thought to utilize 3d orbitals to form M–P pi-bonding, but it is now accepted that d-orbitals on phosphorus are not involved in bonding as they are too high in energy.[13][14]

IUPAC definition of Back Donation

[edit]

The full IUPAC definition of back donation is as follows:

A description of the bonding of π-conjugated ligands to a transition metal which involves a synergic process with donation of electrons from the filled π-orbital or lone electron pair orbital of the ligand into an empty orbital of the metal (donor–acceptor bond), together with release (back donation) of electrons from an nd orbital of the metal (which is of π-symmetry with respect to the metal–ligand axis) into the empty π*-antibonding orbital of the ligand.[15]

See also

[edit]

References

[edit]
  1. ^ Miessler, Gary L.; Tarr, Donald A. (1999). Inorganic chemistry (2nd ed.). Upper Saddle River, N.J: Prentice Hall. p. 338. ISBN 978-0-13-841891-5.
  2. ^ Cotton, Frank Albert; Wilkinson, Geoffrey; Murillo, Carlos A., eds. (1999). Advanced inorganic chemistry (6th ed.). New York: Wiley. ISBN 978-0-471-19957-1.
  3. ^ Housecroft, C. E.; Sharpe, A. G. (2005). Inorganic Chemistry (2nd ed.). Pearson Prentice-Hall. p. 702. ISBN 978-0-130-39913-7.
  4. ^ Crabtree, Robert H. (2014). The Organometallic Chemistry of the Transition Metals (6th ed.). Wiley. pp. 105–106. ISBN 978-1-11813807-6.
  5. ^ a b Elias, Anil J.; Gupta, B D (January 1, 2013). Basic Organometallic Chemistry: Concepts, Syntheses and Applications (2nd ed.). Universities Press. ISBN 978-81-7371-874-8.
  6. ^ a b c Hartwig, John Frederick (2010). Organotransition metal chemistry: from bonding to catalysis. Sausalito (Calif.): University science books. ISBN 978-1-891389-53-5.
  7. ^ Elschenbroich, Christoph (2011). Organometallics (3., compl. rev. and extended ed.). Weinheim: WILEY-VCH. ISBN 978-3-527-29390-2.
  8. ^ Zhao, Haitao; Ariafard, Alireza; Lin, Zhenyang (2006-08-01). "In-depth insight into metal–alkene bonding interactions". Inorganica Chimica Acta. Protagonists in Chemistry: Professor D.M.P. Mingos. 359 (11): 3527–3534. doi:10.1016/j.ica.2005.12.013. ISSN 0020-1693.
  9. ^ Mondal, A.; Breitwieser, K.; Danes, S.; Grünwald, A.; Heinemann, F. W.; Morgenstern, B.; Müller, F.; Haumann, M.; Schütze, M.; Kass, D.; Ray, K.; Munz, D. (2024). "pi-Lewis Base Activation of Carbonyls and Hexafluorobenzene". Angewandte Chemie International Edition e202418738. doi:10.1002/anie.202418738.
  10. ^ Orpen, A. G.; Connelly, N. G. (1990). "Structural systematics: the role of P–A σ* orbitals in metal–phosphorus π-bonding in redox-related pairs of M–PA3 complexes (A = R, Ar, OR; R = alkyl)". Organometallics. 9 (4): 1206–1210. doi:10.1021/om00118a048.
  11. ^ Crabtree, Robert H. (2009). The Organometallic Chemistry of the Transition Metals (5th ed.). Wiley. pp. 99–100. ISBN 978-0-470-25762-3.
  12. ^ Dunne, B. J.; Morris, R. B.; Orpen, A. G. (1991). "Structural systematics. Part 3. Geometry deformations in triphenylphosphine fragments: A test of bonding theories in phosphine complexes". Journal of the Chemical Society, Dalton Transactions: 653. doi:10.1039/dt9910000653.
  13. ^ Gilheany, D. G. (1994). "No d Orbitals but Walsh Diagrams and Maybe Banana Bonds: Chemical Bonding in Phosphines, Phosphine Oxides, and Phosphonium Ylides". Chem. Rev. 94 (5): 1339–1374. Bibcode:1994ChRv...94.1339G. doi:10.1021/cr00029a008. PMID 27704785.
  14. ^ Fey, N.; Orpen, A. G.; Harvey, J. N. (2009). "Building ligand knowledge bases for organometallic chemistry: Computational description of phosphorus(III)-donor ligands and the metal–phosphorus bonds". Coord. Chem. Rev. 253 (5–6): 704–722. doi:10.1016/j.ccr.2008.04.017.
  15. ^ McNaught, A. D.; Wilkinson, A. (2006). IUPAC. Compendium of Chemical Terminology, 2nd ed. (the "Gold Book"). Oxford: Blackwell Scientific Publications. doi:10.1351/goldbook. ISBN 978-0-9678550-9-7.

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Wikipedia

Pi backbonding

In chemistry, pi backbonding or π backbonding is a π-bonding interaction between a filled (or half filled) orbital of a transition metal atom and a vacant orbital on an adjacent ion or molecule. In this type of interaction, electrons from the metal are used to bond to the ligand, which dissipates excess negative charge and stabilizes the metal. It is common in transition metals with low oxidation states that have ligands such as carbon monoxide, olefins, or phosphines. The ligands involved in π backbonding can be broken into three groups: carbonyls and nitrogen analogs, alkenes and alkynes, and phosphines. Compounds where π backbonding is prominent include Ni(CO)4, Zeise's salt, and molybdenum and iron dinitrogen complexes.

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Wikipedia

Pi bond

In chemistry, pi bonds (π bonds) are covalent chemical bonds, in which each of two lobes of an orbital on one atom overlap with two lobes of an orbital on another atom, and in which this overlap occurs laterally. Each of these atomic orbitals has an electron density of zero at a shared nodal plane that passes through the two bonded nuclei. This plane also is a nodal plane for the molecular orbital of the pi bond. Pi bonds can form in double and triple bonds but do not form in single bonds in most cases. The Greek letter π in their name refers to p orbitals, since the orbital symmetry of the pi bond is the same as that of the p orbital when seen down the bond axis. One common form of this sort of bonding involves p orbitals themselves, though d orbitals also engage in pi bonding. This latter mode forms part of the basis for metal-metal multiple bonding.

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Wikipedia

Dewar–Chatt–Duncanson model

The Dewar–Chatt–Duncanson model is a model in organometallic chemistry that explains the chemical bonding in transition metal alkene complexes. The model is named after Michael J. S. Dewar, Joseph Chatt and L. A. Duncanson. The Dewar–Chatt–Duncanson model describes the binding of a transition metal to the C=C bond. The alkene donates electron density into a π-acid metal d-orbital from a C−C π bonding orbital between the carbon atoms, allowing the alkene to behave as a σ-donor to the metal. The metal donates electrons back from a (different) filled d-orbital into the empty π* antibonding orbital, allowing the alkene to function as a π-acceptor. Both of these effects tend to reduce the carbon-carbon bond order, leading to an elongated C−C distance and a lowering of its vibrational frequency. Overall, forward donation from the alkene is more prominent for electronegative, late transition metals (e.g., Cu, Au, Pt, Hg), resulting in complexes electrophilic at the carbon, while backdonation from the metal is more prominent for electropositive, early transition metals, provided the metal is not d0 (i.e., the metal lacks d electrons to donate). For complexes of very early transition metals like (iPrO)2Ti(η2-C2H4)), for example, the backbonding can be so important that the molecule is better described as a metallacyclopropane with metal in the n+2 oxidation state (in this example, Ti(IV)), rather than an alkene π-complex of the metal in oxidation state n (in this example, Ti(II)). These complexes are therefore nucleophilic at the carbon, possessing significant carbanionic character. Progressing across the d-block from left to right, d orbital energies steadily decrease. As a consequence, although the late transition metals often have many d electrons, these electrons tend to reside in orbitals that are too low-lying to allow for efficient backbonding to the π* orbital. In particular, backbonding has been found computationally to contribute less than 20% of the bond energy for d10 group 10 and 11 alkene complexes, and is energetically negligible (

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Wikipedia

Tolman electronic parameter

The Tolman electronic parameter (TEP) is a measure of the electron donating or withdrawing ability of a ligand. It is traditionally determined by measuring the frequency of the A1 C-O vibrational mode (ν(CO)) of a (pseudo)-C3v symmetric complex, [LNi(CO)3] by infrared spectroscopy, where L is the ligand of interest. [LNi(CO)3] was chosen as the model compound because such complexes are readily prepared from tetracarbonylnickel(0). Analogous tetrahedral and square planar complexes, such as rhodium carbonyl chlorides, have also been utilized in measuring the chelating strength of a ligand. The shift in ν(CO) reflects how L alters metal→CO backbonding through its σ‑donor and π‑acceptor (or π‑donor) character. Strong σ‑donor/π‑donor ligands increase metal→CO backbonding, weakening the C≡O bond and lowering ν(CO), whereas weaker donors or π‑acceptors decrease backbonding and raise ν(CO). This balance between donation and back‑donation governs ligand effects on metal–ligand bond strengths, geometries, and reactivity in other complexes, providing a method of categorizing ligands in order. The analysis was introduced by Chadwick A. Tolman.

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TOPIC OF THE DAY

Greenwood / Black Wall Street

Before the 1921 destruction of Tulsa’s Greenwood District, Black residents had created a remarkable center of business and community life. The district included stores, professional offices, entertainment venues and homes owned by Black citizens. Understanding Greenwood means learning what was built—not only what was burned.

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TRIVIA QUESTION OF THE DAY

Which Black woman became the first elected to the United States Congress?

Shirley Chisholm, elected in 1968.