CHM 303 Organometallic Chemistry Past Paper 2022 PDF

CHM 303: ORGANOMETALLIC CHEMISTRY Dr. Segun A. Ogundare Lecture content: Electron rule, bonding. Preparation of Organo-transition metal compounds. Reactions and structures of organometallics. The Organic chemistry of Ferrocene and other related compounds. Metal carbonyl, metal carbonyl halides and...

CHM 303: ORGANOMETALLIC CHEMISTRY Dr. Segun A. Ogundare Lecture content: Electron rule, bonding. Preparation of Organo-transition metal compounds. Reactions and structures of organometallics. The Organic chemistry of Ferrocene and other related compounds. Metal carbonyl, metal carbonyl halides and hydrides. The Effective Atomic Number (EAN) rule or the 18-electron rule and bonding scheme The effective atomic number rule also known as the 18-electron rule states that stable organometallic compounds should have 18 electrons in their outermost (valence) shell. This rule was first proposed by Sidgwick and extended later by Bailey. Therefore, it is also referred to as sidgwick-Bailey's rule. Although there are many exceptions to this rule, it still provides useful guides to the chemistry of many organometallic complexes. The 18-electrons in the valence shell of the metal will occupy ns, np and (n-1)d orbitals a total of 9 orbitals, which can accommodate two electrons each giving a total of 18 electrons for a complete shell. This makes an organometallic complex to be kinetically stable. By this definition, the metal should be a transition metal and n must be greater than 4. If there are less than 18 electrons in the valence shell of the metal, an empty low-lying orbitals into which electrons may be promoted will be available. This will lead to decomposition on slight heating. If there are more than 18 electrons, the excess electron will move to the antibonding orbitals which will reduce the stability by decreasing the bond order due to interaction with bonding electrons. Note that a = a single non-degenerate orbital, e = double degenerate orbitals and t = triple degenerate orbitals. 1 indicates symmetrical to the plane of reflection and 2 indicates unsymmetrical to the plane of reflection. Gerade (g) implies symmetrical to center of inversion and ungerade (u) implies unsymmetrical to center of inversion. Application of 18-electron rule The 18-electron rule helps to predict the stability of organometallic complexes and also helps to predict the most probable structure of newly synthesized organometallic complexes. The application of the 18-electron rule involves the determination of electrons in the valence shell of the metal in zero oxidation state. This is added to the electrons contributed by the ligands and if the complex is positively charged the Lost electron must be removed or deducted from the total electrons. However, if the complex is negatively charged the gained electrons must be added to the total number of electrons. If the compound or complex has 18 electrons in the valence shell of the metal it is considered stable and on heating such complex it will not decompose at relatively low temperature. However, if the number of electrons in the valence shell of a metal is less than or greater than 18, the complex is unstable and may decompose on heating except those that have a special stability associated with 16-electron Square planar complexes. Note that the second row and third row transition metals under these group we'll have the same number of valence electron: nickel, palladium and platinum will have 10 valence electron is zero station state. Similarly, iron osmium and ruthenium will have eight electrons in valence shell in zero oxidation state. The number of electrons donated by the ligands depend on the nature of the ligand and the nature of the bonding below are some examples: Questions i. Using the above given values predict the stability of the following organometallic compounds based on effective atomic number rule. ii. Predict the most probable structures for the complexes given below if the effective atomic number rule is obeyed.1. [(C7H7)Co(CO)3] 2. [Ni(η5-C5H5)(NO)] iii. Draw the most probable structure for [IrCl(PPh3(NO)(CO)]Cl if the number of valence electrons is 16 in the metal. Limitations of the effective atomic number (18-electron) rule There are no known cases in which the meta center in organometallic compounds do not have 18 electrons in their valence shell and they are very stable. Examples include: CH3TiCl3, (CH3)2NbCl3, W(CH3)2, [Rh(CO)2Cl]2, which are: 8, 10, 12 and 16 electron systems respectively. The geometry adopted by these complexes is square planner. Another example is: [Ni(P(cyclohexyl)3)2(C2H4)]. The bulkiness of bonded ligands imposes steric effect on the approach of incoming ligand. this can also impose limitations with regards to the number of ligand that can be accommodated by the metal and limit the number of electrons in the valence shell. Preparation of organo-transition metal compounds The methods of preparing organo-transition metal complexes vary and they depend on the products to be formed. There are several methods only a few will be highlighted in this class with respect to organometallic transition metal complexes. By definition, organometallic compounds are compounds containing at least one metal-to-carbon bond in which the carbon is part of an organic group. The metal-to-carbon bond can be single (M-C), double (M=C) or triple (M≡C) bond. Transition metals can also coordinate with π-bond(s) in hydrocarbons. The organometallic chemistry with regards to each class will be discussed separately. However, simple reactions leading to the formation of these organometallic compounds will be presented here. Preparation of alkyl-transition metal compounds Redox reaction: 2Zn + 2CH3I → Zn(CH3)2 + ZnI2 Double replacement reaction (Metathesis): 3Zn(CH3)2 + WCl6 → 3ZnCl2+ W(CH3)6 ZrCl4 + 4(C6H5-CH2)MgCl → Zr(CH2-C6H5)4 + 4MgCl2 Preparation of carbene-transition metal compounds Preparation of carbyne-transition metal compounds Preparation of carbonyl-transition metal compounds Direct reaction: Ni + 4CO → Ni(CO)4 Fe + 5CO → Fe(CO)5 2Co + 8CO → Co2(CO)8 Reaction with metal salt 2CrCl3 + 12CO + 3Zn/Hg→ 2Cr(CO)6 + 3ZnCl2 + 3Hg (Zn is the reducing agent and Cr3+ is reduced to Cr) VCl3 + 6CO + 3Na → V(CO)6 + 3NaCl (Na is the reducing agent and V3+ is reduced to V) Reactions organometallics The reactivity of organometallics depends on the nature of the organic ligands and the metal to which they are attached. The reactions of organometallics are numerous and few with important applications will be highlighted. 1. Ligand substitution reaction Fe(CO)5 + P(C6H5)3 → (CO)4FeP(C6H5)3 + CO 2. Insertion reaction 3. Sigma-bond metathesis (double replacement reaction) This reaction involves the formation of a C—H bond in methane as by-product. This occurs as a C—H bond of benzene is cleaved, but the oxidation state of the scandium center remains (+3). This is called sigma-bond metathesis. 4. Olefin Metathesis The olefin metathesis reaction (the subject of 2005 Nobel Prize in Chemistry) can be thought of as a reaction in which all the carbon-carbon double bonds in an olefin (alkene) are cut and then rearranged in a statistical fashion: If one of the product alkenes is volatile (such as ethylene) or easily removed, then the reaction shown above can be driven completely to the right. Likewise, using a high pressure of ethylene, internal olefins can be converted to terminal olefins. There are a wide variety of variants on this reaction as is discussed below. Mechanism The commonly accepted mechanism for the olefin metathesis reaction was proposed by Chauvin and involves a [2+2] cycloaddition reaction between a transition metal alkylidene complex and the olefin to form an intermediate metallacyclobutane. This metallacycle then breaks up in the opposite fashion to afford a new alkylidene and new olefin. If this process is repeated enough, eventually an equilibrium mixture of olefins will be obtained. The organic chemistry of Ferrocene and other related compounds In 1951, an attempt to synthesize fulvalene from cyclopentadienyl bromide, Grignard reagent ((C5H5)MgBr) was reacted with FeCl3. This reaction produced an orange solid having the formula (C5H5)2Fe, ferrocene instead of fulvalene. It was noted that all the carbon atoms in the cyclopentadienyl rings were equidistant from the metal ion, making the compounds appeared like a sandwich. There are other related organometallic compounds with sandwich structures and they are generally called metallocenes. Most metallocenes are prepared by reaction of the corresponding metal salt with NaC5H5. Other metallocenes have similar structures but do not necessarily obey the 18-electron rule. For example, chromocene, cobaltocene and nickelocene are structurally similar 16-, 19- and 20-electron species. The metallocenes are generally unstable in air with exception of Ferrocene, Ruthenocene and Osmocene which obey the 18-electron rule. Reactions of metallocenes Metallocenes are susceptible to electrophilic attacks at the cyclopentadienyl rings. An example is formation of acetyl derivative using acetyl chloride. A double substitution can also occur on further acetylation. Similar reaction will occur in the presence of acetic anhydride. Ferrocene can also be converted to ferrocenic acid by formylation and subsequent oxidation of the intermediate. Explain why the Fe-C distance lengthens for [Cp2Fe]+, while the Co-C distance shortens for [Cp2Co]+. Note the various trends in the bond distances. The changes in the neutral Fe, Co, Ni metallocenes are a direct result of going from 18e-(Fe) to 19e-(Co) to 20e-(Ni) counts. The extra electrons for the Co and Ni complexes are going into M-Cp antibonding orbitals, which are delocalized and progressively weaken the M-Cp bonding, leading to the increase in bond distances. Metal carbonyl, metal carbonyl halides and hydrides CO is considered as a suitable ligand in organometallic chemistry because it functions as a Lewis base contributing two electrons to metal orbitals. Although CO is inorganic molecule but the large number of compounds formed by its reaction with transition metals makes it important to include this class of compounds in the study of organometallic chemistry. There are binary metal carbonyls containing one metal examples: Fe(CO)5, Ni(CO)4. These are called primary metal carbonyl. With two or more metals the carbonyls are called secondary metal carbonyls examples; Fe2(CO)8, Os3(CO)12. When halogen is also coordinated to the metal, the compounds are called metal carbonyl halides examples; Mn(CO)5I, Co(CO)4Iand with hydrogen coordinated to the metal, the compounds are called metal carbonyl hydrides examples: HMn(CO)5, HCo(CO)4 and H2Fe(CO)4. Similarly, if NO is coordinated to the metal, the compounds are called metal carbonyl nitrosyls examples; Co(CO)3NO and MnCO(NO)3. Preparation methods of metal carbonyl Direct reaction: Ni + 4CO → Ni(CO)4 Fe + 5CO → Fe(CO)5 2Co + 8CO → Co2(CO)8 Photolysis of primary metal carbonyl to form secondary metal carbonyl 2Fe(CO)5 → Fe2(CO)9 + CO (UV radiation at room temperature and the solvent is n-hexane.) 2Os(CO)5 → Os2(CO)9 + CO (UV radiation at -40ºC and the solvent is n-hexane.) Reaction with metal salt 2CrCl3 + 12CO + 3Zn/Hg→ 2Cr(CO)6 + 3ZnCl2 + 3Hg (Zn is the reducing agent and Cr3+ is reduced to Cr) VCl3 + 6CO + 3Na → V(CO)6 + 3NaCl (Na is the reducing agent and V3+ is reduced to V) Reduction of secondary binary metal carbonyl to form primary metal carbonyl Mn2(CO)10 + 2Na/Hg → 2Na+ [Mn(CO)5]- + 2Hg (at room temp. and the solvent is THF.) Re2(CO)10 + 2Na/Hg → 2Na+ [Re(CO)5]- + 2Hg (at 40ºC and the solvent is THF.) Other metal carbonyls 2Na+ [Re(CO)5]- + 2I2 → 2[Re(CO)5I] + 2NaI Co2(CO)8 + 2NO → 2Co(CO)3NO + 2CO Co2(CO)8 + H2 → 2HCo(CO)4 Mn2(CO)10 + H2 → 2HMn(CO)5 Bonding in metal carbonyl The bonding in metal carbonyls can be described from the perspective of molecular orbital diagram of CO, which is presented below. The p-orbitals in O and C overlap to give two π-bonds and a σ- bond. The highest energy occupied molecular orbital (HOMO) is σ-bonding orbital which is used to coordinate to the metal. However, the presence of unoccupied empty π*-antibonding orbitals makes electrons from the d-orbitals of the metal to donate electron to the CO orbitals. This is called back bonding or π-bonding scheme. Both the π-bond and σ-bond constitute the bonding mode in metal carbonyls. The HOMO has its largest lobe on carbon. It is through this orbital, occupied by an electron pair, that CO exerts its σ-donor function, donating electron density directly toward a suitable metal orbital, which could be an empty d or hybrid orbital. CO also has two empty π*-antibonding orbitals (the lowest unoccupied molecular orbital (LUMO); these also have larger lobes on carbon than on oxygen. A metal atom having electrons in a d-orbital of suitable symmetry can donate electron density to these π*-antibonding orbitals. These σ-donor and π-acceptor interactions are illustrated below; The σ-bonding interaction increases the electron density on the metal and decreases the electron density on the CO ligand. The π-bonding interaction decreases the electron density on the metal and increases the electron density on the CO ligand. Both effects ‘reinforce’ each other. The bond strength is dependent of a number of factors which including the charge on the complex and presence of other ligands in the environment of the metal. The bonding scheme proposed can be supported by experimental evidence. The evidences are from infrared spectroscopic and X-ray crystallographic studies. The IR will show the change in bond stretching vibration of CO due to coordination with metal. This will lead to reduction in the energy of vibration as the bond in CO becomes weaker. The C-O stretch in organometallic complexes is often very intense (stretching the C-O bond results in a substantial change in dipole moment), and its energy often provides valuable information about the molecular structure. Free CO (uncoordinated) has a C-O stretch at (2143 cm–1). Cr(CO)6, on the other hand, has its C-O stretch at (2000 cm–1). The lower energy for the stretching mode means that the C-O bond is weaker in Cr(CO)6. Both σ donation and π acceptance would be expected to weaken the CO bond and to decrease the energy necessary to stretch that bond. X-ray crystallography also provide additional information on the bonding in metal carbonyls. The C-O distance has been measured as 112.8 pm in uncoordinated CO. Weakening of the C-O bond would be expected to cause this distance to increase. Such an increase in bond length is found in complexes containing CO, with C-O distances approximately 115 pm for many carbonyls. The charge on a carbonyl complex is also reflected in its infrared spectrum. Five isoelectronic hexacarbonyls have the following C-O stretching bands (compare with v(CO) = 2143 cm-1 for free CO): The positions of the C-O stretching vibrations can also be affected by solvent when the IR is measured in solution. Similarly, counter ions can also influence the position of IR vibrations of metal carbonyl. Hence the position can vary with change in counter ion and also the position of the vibration band can also vary with different solvent and when compared to neat (without solvent) samples. Note that the higher the negative charge on the complex the weaker the CO bond length. [Ti(CO)6]2- contains the most highly reduced metal, formally containing Ti2-; this means that titanium has the weakest ability to attract electrons and the greatest tendency to back donate electron density to CO. The formal charges on the metals increase from (-2) for [Ti(CO)6]2- to (+2) for [Fe(CO)6]2+ with corresponding increase in the energy of vibration. The titanium in [Ti(CO)6]2-, with the most negative formal charge, has the strongest tendency to donate to CO. The consequence is strong population of the π* orbitals of CO in [Ti(CO)6]2- and reduction of the strength of the C-O bond. In general, the more negative the charge on the organometallic species, the greater the tendency of the metal to donate electrons to the π* orbitals of CO, and the lower the energy of the C-O stretching vibrations. Cationic carbonyl complexes such as [Fe(CO)6]2+ have C-O stretching bands even higher in energy than those in free CO. The CO ligand in the complex does not engage in significant π-acceptor activity therefore weakening of the C-O bond through this interaction should be minimal. It is important to note that the positive charge of the metal will cause the distortion of the electron cloud of the CO leading to polarization of the CO bond. In free CO, the electrons are polarized toward the more electronegative oxygen. For example, the electrons in the p orbitals are concentrated nearer to the oxygen atom than to the carbon. The presence of a transition metal cation reduces the polarization in the C-O bond by attracting the bonding electrons. The consequence is that the electrons in the positively charged complex are more equally shared by the carbon and the oxygen, giving rise to a stronger bond and a higher- energy C-O stretch. The very high ν(CO) bands result from weak back donation. When the frequency of carbonyls appears at higher energy band of free CO, the complexes are sometimes called non-classical carbonyls. Bridging Carbonyls When CO is coordinated to more than one metal atom in a complex, the CO is said to be bridging. The bridging mode is supported by the position of the C-O stretching vibration band. When the CO is bonded to two metal atoms, the CO IR stretching vibration energy is reduced and further decrease is noted when the CO is coordinated to three or more metal atoms. This is because the electron density will increase in the bridging CO as more metal atoms contribute more electrons to the π*-antibonding orbital of the CO. Ordinarily, terminal and bridging carbonyl ligands can be considered 2 electron donors, with the donated electrons shared by the metal atoms in the bridging modes. The bridging CO is a 2-electron donor overall, with a single electron donated to each metal in doubly bridging mode. Properties of metal carbonyl 1. They are mostly crystalline solids at room temperature with few exceptions such as Fe(CO)5, Ru(CO)5 and Os(CO)5 2. They are generally soluble in organic solvents and highly toxic if ingested 3. They are stable to air when the EAN rule is obeyed but becomes less stable if the number of electron is more or less than 18 with exception of few. Those that are unstable are stored under nitrogen or argon. 4. Metals with odd atomic number cannot obey the EAN rule. By simple addition of CO ligand, since the resultant moiety will have an odd number of electrons. In such case there are several option open to these metals by which the EAN rule can be obeyed a. The addition of an electron by reducing agent to form an anion such as [V(CO)6]-.  b. The electron deficient moiety can bond covalently with an atom or group that also has single unpaired electron available, example; hydrogen or chlorine: HM(CO)n or M(CO)nCl. c. If no either species are available with which to react, two moieties each with an odd atom can form a dimer with resultant pairing of the odd electrons, examples Questions 1. Determine if EAN rule is obeyed in the following complexes and where the rule is not obeyed suggest possible reactions to attain stability. i. V(CO)6, ii. Cr(CO)6 and iii. Mn(CO)6 iv. Fe(CO)4, v. Co(CO)4 and vi. Ni(CO)4 2. Briefly discuss the variation observed in the IR stretching vibrational bands observed in the following; i. [Fe(CO)4]2- (1790 cm-1) ii. [Co(CO)4]- (1890 cm-1) and iii. Ni(CO)4 (2060 cm-1) and [Cu(CO)4]+ (2180 cm-1) 3. With simple illustration describe the nature of bonding in metal carbonyls and state two experimental evidences that justify this description.