organo metallic

The 18 electron rule

Just as organic chemists have their octet rule for organic compounds, so do organometallic chemists have the 18 electron rule. And just as the octet rule is often violated, so is the 18 electron rule. However, both serve a useful purpose in predicting reactivity. Each derives from a simple count of the number of electrons that may be accommodated by the available valence orbitals (one s and three p for organic chemists; organometallic chemists get five bonus d-orbitals in which to place their electrons).

Counting electrons in organometallic complexes

Knowing how many valence electrons "belong to" a transition metal complex allows us to make predictions about the mechanisms of reactions and the possible modes of reactivity. There are two distinct methods that are used to count electrons, theneutral or covalent methodand the effective atomic number orionic method. While this may seem confusing, these are simply two different accounting systems that give us the same final answer. Although having two systems may seem confusing, it at least provides us a convenient way to double check our answer.

What are d-electrons, anyway?

While we teach our students in freshman chemistry that the periodic table is filled in the order [Ar]4s23d10, this turns out to be true only for isolated metal atoms. When we put a metal ion into an electronic field (surround it with ligands), the d-orbitals drop in energy and fill first. Therefore, the organometallic chemist considers the transition metal valence electrons toallbe d-electrons. There are certain cases where the 4s23dxorder does occur, but we can neglect these in our first approximation.

Therefore, when we ask for the d-electron count on a transition metal such as Ti in the zero oxidation state, we call it d4, not d2. For zero-valent metals, we see that the electron count simply corresponds to the column it occupies in the periodic table. Hence, Fe is in the eighth column and is d8(not d6) and Re3+is d4(seventh column for Re, and then add 3 positive charges...or subtract three negative ones). Now that we can assign a d-electron count to a metal center, we are ready to determine the electronic contribution of the surrounding ligands and come up with our overall electron count.

Method 1: The ionic (charged) model

The basic premise of this method is that we remove all of the ligands from the metal and, if necessary, add the proper number of electrons to each ligand to bring it to a closed valence shell state. For example, if we remove ammonia from our metal complex, NH3has a completed octet and acts as a neutral molecule. When it bonds to the metal center it does so through its lone pair (in a classic Lewis acid-base sense) and there is no need to change the oxidation state of the metal to balance charge. We call ammonia a neutral two-electron donor.

In contrast, if we remove a methyl group from the metal and complete its octet, then we formally have CH3-. If we bond this methyl anion to the metal, the lone pair forms our metal-carbon bond and the methyl group acts as a two-electron donor ligand. Notice that to keep charge neutrality we must oxidize the metal by one electron (i.e. assign a positive charge to the metal). This, in turn, reduces the d-electron count of the metal center by one. We'll seeseveral examplesbelow.

Method 2: The covalent (neutral) model

The major premise of this method is that we remove all of the ligands from the metal, but rather than take them to a closed shell state, we do whatever is necessary to make them neutral. Let's consider ammonia once again. When we remove it from the metal, it is a neutral molecule with one lone pair of electrons. Therefore, as with the ionic model, ammonia is a neutral two electron donor.

But we diverge from the ionic model when we consider a ligand such as methyl. When we remove it from the metal and make the methyl fragment neutral, we have a neutral methylradical. Both the metal and the methyl radical must donate one electron each to form our metal-ligand bond. Therefore, the methyl group is a one electron donor, not a two electron donor as it is under the ionic formalism. Where did the other electron "go"? It remains on the metal and is counted there. In the covalent method, metals retain their full complement of d electrons because we never change the oxidation state from zero; i.e. Fe will always count for 8 electrons regardless of the oxidation state and Ti will always count for four.

Notice that this method does not give us any immediate information about the formal oxidation state of the metal, so we must go back and assign that in a separate step. For this reason, many chemists (particularly those that work with high oxidation state complexes) prefer the ionic method.

Electron donation of common ligands

In the table below are some common transition metal ligands and the number of electrons that each donates to a metal center. Some ligands can donate a variable number of electrons. For example, an alkoxide, M-OR, can donate two to six electrons depending on the hybridization of the oxygen atom.

The two methods compared: some examples

The most critical point we should remember is thatlike oxidation state assignments, electron counting is a formalism and does not necessarily reflect the distribution of electrons in the molecule. However, these formalisms are very useful to us, and both will give us the same final answer.

Consider the following simple examples. Notice how some ligands donate the same number of electrons no matter which formalism we choose, while the number of d-electrons and donation of the other ligands can differ. All we have to do isremember to be consistentand it will work out for us.

Self-Test

In each of the following examples, select the oxidation state of the transition metal atom, dncount (assign it using theionic model), and number of valence electrons at each metal center. Then hit the Submit Answer button for that compound to see if you are correct.

If you do not yet know your periodic table, you can either visit thePeriodic Table Challengeto test your skill or you can view a complete periodic table inthis popup window.

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Oxidation State:

dncount:

Valence Electrons:

Thishalf-sandwichis used as a gasoline anti-knock additive.

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Oxidation State:

dncount:

Valence Electrons:

Free hint: This is quite similar to an example shown above.

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Oxidation State:

dncount:

Valence Electrons:

Answer for either equivalent Re, not both. This was the first example of a metal-metal quadruple bond.


Oxidation State:

dncount:

Valence Electrons:

This was the first example of aSchrock alkylidene complex.

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Oxidation State:

dncount:

Valence Electrons:

Notice that isnotacyclopentadienyl ligand.


Oxidation State:

dncount:

Valence Electrons:

A classic example of ametallocene.

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Oxidation State:

dncount:

Valence Electrons:

Give your answer for either equivalent iridium center.


Oxidation State:

dncount:

Valence Electrons:

This has two eachhydride,dihydrogenandphosphineligands.

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Oxidation State:

dncount:

Valence Electrons:

Hint 1: the sulfur ligands are called dithiocarbamates (dtc). Hint 2: look closely at theimido ligands.


Oxidation State:

dncount:

Valence Electrons:

How are you going to count the alkoxide ligands? Hint: see thedonation chartabove.

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Now, we will learning about the 18-electron rules in organometallic chemistry. Seperti biasa saya merujuk dari buku miessler dan huheey dan materi kuliah. Secara sederhana dan basic suatu molekul kimia dikatakan stabil apabila telah memiliki jumlah electron menyerupai gas mulia. Aturan ini kita kenal dengan aturan octet. Kemudian kenapa pada kimia organologam dikenal 18-elektron dan pada molekul sederhana hanya 8-elektron? Secara sederhana dapat kita analogkan 18-elektron dengan 8-elektron, dan analog itu begini: aturan 8-elektron digunakan untuk senyawaan golongan utama, yang artinya merepresentasikan jumlah electron pada kulit valensi terisi penuh (S2P6), sedangkan 18-elektron berhubungan dengan jumlah electron valensi untuk logam transisi (s2p6d10). Dari penjelasan tersebut dapat diketahui mengapa organologam menggunakan aturan 18-elektron, karena organologam merupakan senyawaan yang mengandung atom pusat logam transisi (sebagian besar) dan senyawa organic (hidrokarbon).

Oke setelah tahu penjelasan singkat tentang aturan 18-elektron, sekaran mari menginjak ke bagaimana cara menghitung jumlah electron senyawa organologam dengan system 18-elektron. Aturan 18-elektron ini terbagi menjadi dua metode, pertama adalah metodeDonor Pairdan yang kedua metodeNeutral Ligand.

1. Donor Pair Method (Method A)

Pada metode ini melibatkan ligan sebagai pendonor pasangan electron ke logam. Untuk menentukan jumlah total electron, kita harus menghitung juga muatan setiap ligand dan menentukan tingkat atau bilangan oksidasi formal dari atom pusat.

Contoh: dikarbonil kloro pentahapto siklopentadienil besi (II) (5-C5H5)Fe(CO)2Cl

5-C5H5-mendonorkan 3 pasang e-, CO mendonorkan 2 elektron (karena terdapat 2 CO maka dikali 2), Fe(II) = [Ar]4s03d6jadi mendonorkan 6 elektron, sehingga jika dijumlahkan:

Fe (II) : 6 e-

5-C5H5: 6 e-

2(CO) : 4 e-

Cl- : 2 e-

Total eletron : 18 elektron

2. Neutral Ligand Method (Method B)

Pada metode ini kita akan menggunakan jumlah electron yang akan didonasikan oleh ligand tetapi dalam keadaan netral. Pada ligan anorganik sederhana, jumlah electron yang didonasikan sama dengan muatan negatifnya sebagai ion bebeas. Misalnya Cl donor 1 e- (muatan ion bebas -1), O donor 2 e- (muatan ion bebeas -2), N donor 3 e- (muatan ion bebas -3). Pada metode ini kita tidak memerlukan penentuan bilangan oksidasi dari atom pusat.

Contoh: dikarbonil kloro pentahapto siklopentadienil besi (II) (5-C5H5)Fe(CO)2Cl

Fe mendonorkan 8 e- sesuai konfigurasi elektronnya, 5-C5H5mendonorkan 5 e- (kita mempertimbangan ligan ini sebagai ligan netral), CO mendonorkan 2 e-, dan Cl mendonorkan 1 e- (sebagai spesies netral)

Fe : 8 e-

5-C5H5 : 5 e-

2(CO) : 4 e-

Cl : 1 e-

Total = 18 e-

18-Electron rule

From Wikipedia, the free encyclopedia

The18-electron ruleis a rule used primarily for predicting formulae for stable metal complexes.[1]The rule is based on the fact that thevalence shellsoftransition metalsconsist of nine valence orbitals, which collectively can accommodate 18electronsas either bonding or nonbonding electron pairs. This means that, the combination of these nineatomic orbitalswithligandorbitals creates ninemolecular orbitalsthat are either metal-ligand bonding or non-bonding. When a metal complex has 18 valence electrons, it is said to have achieved the same electron configuration as thenoble gasin the period. The rule and its exceptions are similar to the application of theoctet ruleto main group elements. The rule is not helpful for complexes of metals that are not transition metals, andinteresting or useful transition metal complexes will violate the rulebecause of the consequences deviating from the rule bears on reactivity. The rule was first proposed by American chemistIrving Langmuirin 1921.[1]

HYPERLINK "http://en.wikipedia.org/wiki/18-Electron_rule" \l "cite_note-2" [2]Contents

[hide] 1Applicability of the 18-electron rule 1.1Consequences for reactivity 1.2Alternative analysis 2Exceptions to the 18-electron rule 2.116electron rule 2.2Bulky ligands 2.3High-spin complexes 2.4Pi-donating ligands 2.5Combinations of effects 2.6Higher electron counts 3See also 4References 5Further readingApplicability of the 18-electron rule[edit]The rule usefully predicts the formulae for low-spin complexes of the Cr, Mn, Fe, and Co triads. Well-known examples includeferrocene,iron pentacarbonyl,chromium carbonyl, andnickel carbonyl.

Ligands in a complex determine the applicability of the 18-electron rule. In general, complexes that obey the rule are composed at least partly ofpi-acceptor ligands(also known as -acids). This kind of ligand exerts a very strongligand field, which lowers the energies of the resultant molecular orbitals and thus favorably occupied. Typical ligands includeolefins,phosphines, andCO. Complexes of -acids typically feature metal in a low-oxidation state. The relationship between oxidation state and the nature of the ligands is rationalized within the framework of backbonding.

Consequences for reactivity[edit]Compounds that obey the 18 VE rule are typically "exchange inert." Examples include[Co(NH3)5Cl]2+,Mo(CO)6, and[Fe(CN)6]4. In such cases, in general ligand exchange occurs viadissociative substitutionmechanisms, wherein the rate of reaction is determined by the rate of dissociation of a ligand. On the other hand, 18-electron compounds can be highly reactive toward electrophiles such as protons, and such reactions are associative in mechanism, being acid-base reactions.

Complexes with fewer than 18 valence electrons tend to show enhanced reactivity. Thus, the 18-electron rule is often a recipe for non-reactivity in either astoichiometricor acatalyticsense.

Alternative analysis[edit]In the prevalentligand fieldanalysis, the valence p orbitals on the metal participate in metal-ligand bonding, albeit weakly. Some new theoretical treatments do not count the metal p-orbitals in metal-ligand bonding,[3]although these orbitals are still included aspolarization functions. This results in a duodectet[3](12) rule which accommodates all low-spin complexes including linear 14e complexes such asTollen's reagentand square planar 16e complexes as well as implies that such transition metal complexes arehypervalent, but has yet to be adopted by the general chemistry community.

Exceptions to the 18-electron rule[edit]-donor or -donor ligands with small interactions with the metal orbitals lead to a weakligand fieldwhich increases the energies of t2gorbitals. Thesemolecular orbitalsbecomenon-bondingor weakly anti-bonding orbitals (small oct). Therefore, addition or removal of electron has little effect on complex stability. In this case, there is no restriction on the number of d-electrons and complexes with 12 -22 electrons are possible. Small octmakes filling eg* possible ( > 18e-) and -donor ligands can make t2gantibonding ( < 18 e-). These types of ligand are located in low to medium of thespectrochemicalseries. For example: [TiF6]2(Ti4+, d0, 12 e), [Co(NH3)6]3+(Co3+, d6, 18 e), [Cu(OH2)6]2+(Cu2+, d9, 21 e) In tems of metal ions, octincreases down a group as well as increasingoxidation number. Strong ligand fields lead tolow-spincomplexes which cause some exceptions to 18-electron rule.

16electron rule[edit]A popular class of complexes that violate the 18e rule are the 16e complexes with d8configurations. Allhigh-spind8metal ions areoctahedral(ortetrahedral), but thelow-spind8metal ions are all square planar (Jahn-Tellerdistortion). Important examples of square-planar low-spin d8metal Ions are Ni(II), Pd(II), and Pt(II). At picture below is shown the splitting of the d sub-shell in low-spin square-planar complexes. Examples are especially prevalent for derivatives of the cobalt and nickel triads. Such compounds are typicallysquare-planar. The most famous example isVaska's complex(IrCl(CO)(PPh3)2), [PtCl4]2, andZeise's salt[PtCl3(2-C2H4)]. In such complexes, the dz2orbital is doubly occupied and nonbonding.

Manycatalytic cyclesoperate via complexes that alternate between 18e and square-planar 16 configurations. Examples includeMonsanto acetic acid synthesis,hydrogenations,hydroformylations, olefin isomerizations, and some alkene polymerizations.

Other violations can be classified according to the kinds of ligands on the metal center.

Bulky ligands[edit]Bulky ligands can preclude the approach of the full complement of ligands that would allow the metal to achieve the 18 electron configuration. Examples:

Ti(neopentyl)4(8 VE)

Cp*2Ti(C2H4) (16 VE)

V(CO)6(17 VE)

Cp*Cr(CO)3(17 VE)

Pt(PtBu3)2(14 VE)

Co(norbornyl)4(13 VE)

[FeCp2]+(17 VE)

Sometimes such complexes engage in agostic interactions with the hydrocarbon framework of the bulky ligand. For example:

W(CO)3[P(C6H11)3]2has 16 VE but has a short bonding contact between one C-H bond and the W center.

Cp(PMe3)V(CHCMe3) (14 VE, diamagnetic) has a short V-H bond with the 'alkylidene-H', so the description of the compound is somewhere between Cp(PMe3)V(CHCMe3) and Cp(PMe3)V(H)(CCMe3).

High-spin complexes[edit]High-spin metal complexes have singly occupied orbitals and may not have any empty orbitals into which ligands could donate electron density. In general, there are few or no -acidic ligands in the complex. These singly occupied orbitals can combine with the singly occupied orbitals of radical ligands (e.g.,oxygen), or addition of astrong fieldligand can cause electron-pairing, thus creating a vacant orbital that it can donate into. Examples:

CrCl3(THF)3(15 VE)

[Mn(H2O)6]2+(17 VE)

[Cu(H2O)6]2+(21 VE, see comments below)

Complexes containing strongly pi-donating ligands often violate the 18-electron rule. These ligands include fluoride (F), oxide (O2), nitride (N3), alkoxide (RO), and imide (oxide (RN2). Examples:

[CrO4]2(16 VE)

Mo(=NR)2Cl2(12 VE)

In the latter case, there is substantial donation of the nitrogen lone pairs to the Mo (so the compound could also be described as a 16 VE compound). This can be seen from the short Mo-N bond length, and from the angle Mo - N - C(R), which is nearly 180. Counter-examples:

trans-WO2(Me2PCH2CH2PMe2)2(18 VE)

Cp*ReO3(18 VE)

In these cases, the M=O bonds are "pure" double bonds (i.e., no donation of the lone pairs of the oxygen to the metal), as reflected in the relatively long bond distances.

Pi-donating ligands[edit]Ligands where the coordinating atom bear nonbonding lone pairs often stabilize unsaturated complexes. Metal amides and alkoxides often violate the 18e rule.

Combinations of effects[edit]The above factors can sometimes combine. Examples include

Cp*VOCl2(14 VE)

TiCl4(8 VE)

Higher electron counts[edit]Some complexes have more than 18 electrons. Examples:

Cobaltocene(19 VE)

Nickelocene(20 VE)

The hexaaqua copper(II) ion [Cu(H2O)6]2+(21 VE)

Often, cases where complexes have more than 18 valence electrons are attributed to electrostatic forces - the metal attracts ligands to itself to try to counterbalance its positive charge, and the number of electrons it ends up with is unimportant. In the case of the metallocenes, the chelating nature of the cyclopentadienyl ligand stabilizes its bonding to the metal. Somewhat satisfying are the two following observations: (i) cobaltocene is a strong electron donor, readily forming the 18-electron cobaltocenium cation and (ii) nickelocene tends to react with substrates to give 18-electron complexes, e.g. CpNiCl(PR3) and free CpH.

In the case of nickelocene, the extra two electrons are in orbitals which are weakly metal-carbon antibonding, this is why it often participates in reactions where the M-C bonds are broken and the electron count of the metal will change to 18.[4]Electron countingis a formalism used for classifying compounds and for explaining or predicting electronic structure andbonding.[1]Many rules in chemistry rely on electron-counting:

Octet ruleis used withLewis structuresfor main group elements, especially the lighter ones such ascarbon,nitrogen, andoxygen,

Eighteen electron ruleininorganic chemistryandorganometallic chemistryoftransition metals,

Polyhedral skeletal electron pair theoryforcluster compounds, including transition metals and main group elements such asboronincludingWade'srules forpolyhedralcluster compounds, including transition metals and main group elements and mixtures thereof.

Atoms that do not obey their rule are called "electron-deficient" when they have too few electrons to achieve anoble gasconfiguration, or "hypervalent" when they have too many electrons. Since these compounds tend to be more reactive than compounds that obey their rule, electron counting is an important tool for identifying the reactivity of molecules.

Contents

[hide] 1Counting rules 1.1Neutral counting 1.2Ionic counting 2Electrons donated by common fragments 2.1"Special cases" 3Examples of electron counting 4See also 5ReferencesCounting rules[edit]Two methods of electron counting are popular and both give the same result.

The neutral counting approach assumes the molecule or fragment being studied consists of purely covalent bonds. It was popularized byM.L.H. Greenalong with the L and X ligand notation.[2]

HYPERLINK "http://en.wikipedia.org/wiki/Electron_counting" \l "cite_note-3" [3]It is usually considered easier especially for low-valent transition metals.[citation needed] The "ionic counting" approach assumes purely ionic bonds between atoms. It rewards the user with a knowledge ofoxidation states, which can be valuable.[peacockterm]One can check one's calculation by employing both approaches, though it is important to be aware that most chemical species exist between the purely covalent and ionic extremes.

Neutral counting[edit] This method begins with locating the central atom on the periodic table and determining the number of its valence electrons. One counts valence electrons for main group elements differently from transition metals.

E.g. in period 2: B, C, N, O, and F have 3, 4, 5, 6, and 7 valence electrons, respectively.

E.g. in period 4: K, Ca, Sc, Ti, V, Cr, Fe, Ni have 1, 2, 3, 4, 5, 6, 8, 10 valence electrons respectively.

One is added for everyhalideor other anionic ligand which binds to the central atom through a sigma bond.

Two is added for every lone pair bonding to the metal (e.g. each Lewis base binds with a lone pair). Unsaturated hydrocarbons such as alkenes and alkynes are considered Lewis bases. Similarly Lewis and Bronsted acids (protons) contribute nothing.

One is added for each homoelement bond.

One is added for each negative charge, and one is subtracted for each positive charge.

Ionic counting[edit] This method begins by calculating the number of electrons of the element, assuming an oxidation state

e.g. for a Fe2+has 6 electrons

S2has 8 electrons

Two is added for everyhalideor other anionic ligand which binds to the metal through a sigma bond.

Two is added for every lone pair bonding to the metal (e.g. each phosphine ligand can bind with a lone pair). Similarly Lewis and Bronsted acids (protons) contribute nothing.

For unsaturated ligands such as alkenes, one electron is added for each carbon atom binding to the metal.

Electrons donated by common fragments[edit]LigandElectrons contributed(neutral counting)Electrons contributed(ionic counting)

X12 (X; X = F, Cl, Br, I)

H12 (H)

H10 (H+)

O24(O2)

N36 (N3)

NR322 (NR3; R = H, alkyl, aryl)

CR224 (CR22)

Ethylene22 (C2H4)

cyclopentadienyl56(C5H5)

benzene66 (C6H6)

"Special cases"[edit]The numbers of electrons "donated" by some ligands depends on the geometry of the metal-ligand ensemble. Perhaps the most famous example of this complication is the M-NOentity .[peacockterm]When this grouping is linear, the NO ligand is considered to be a three-electron ligand. When the M-NO subunit is strongly bent at N, the NO is treated as a pseudohalide and is thus a one electron (in the neutral counting approach). The situation is not very different from the -3 vs. -1 allyl. Another unusual ligand from the electron counting perspective is sulfur dioxide.

Examples of electron counting[edit] CH4, for the central C

neutral counting: C contributes 4 electrons, each H radical contributes one each: 4+4(1) = 8 valence electrons

ionic counting: C4contributes 8 electrons, each proton contributes 0 each: 8 + 4(0) = 8 electrons.

Similar for H:

neutral counting: H contributes 1 electron, the C contributes 1 electron (the other 3 electrons of C are for the other 3 hydrogens in the molecule): 1 + 1(1) = 2 valence electrons.

ionic counting: H contributes 0 electrons (H+), C4contributes 2 electrons (per H), 0 + 1(2) = 2 valence electrons

conclusion: Methane follows the octet-rule for carbon, and the duet rule for hydrogen, and hence is expected to be a stable molecule (as we see from daily life)

H2S, for the central S

neutral counting: S contributes 6 electrons, each hydrogen radical contributes one each: 6+2(1) = 8 valence electrons

ionic counting: S2contributes 8 electrons, each proton contributes 0: 8+2(0) = 8 valence electrons

conclusion: with an octet electron count (on sulfur), we can anticipate that H2S would be pseudotetrahedral if one considers the two lone pairs.

SCl2, for the central S

neutral counting: S contributes 6 electrons, each chlorine radical contributes one each: 6+2(1) = 8 valence electrons

ionic counting: S2+contributes 4 electrons, each chloride anion contributes 2: 4+2(2) = 8 valence electrons

conclusion: see discussion for H2S above. Notice that both SCl2and H2S follow the octet rule - the behavior of these molecules is however quite different.

SF6, for the central S

neutral counting: S contributes 6 electrons, each fluorine radical contributes one each: 6+6(1) = 12 valence electrons

ionic counting: S6+contributes 0 electrons, each fluoride anion contributes 2: 0+6(2) = 12 valence electrons

conclusion: ionic counting indicates a molecule lacking lone pairs of electrons, therefore its structure will be octahedral, as predicted byVSEPR. One might conclude that this molecule would be highly reactive - but the opposite is true: SF6is inert, and it is widely used in industry because of this property.

TiCl4, for the central Ti

neutral counting: Ti contributes 4 electrons, each chlorine radical contributes one each: 4+4(1) = 8 valence electrons

ionic counting: Ti4+contributes 0 electrons, each chloride anion contributes two each: 0+4(2) = 8 valence electrons

conclusion: Having only 8e (vs.18possible), we can anticipate that TiCl4will be a good Lewis acid. Indeed, it reacts (in some cases violently) with water, alcohols, ethers, amines.

Fe(CO)5neutral counting: Fe contributes 8 electrons, each CO contributes 2 each: 8 + 2(5) = 18 valence electrons

ionic counting: Fe(0) contributes 8 electrons, each CO contributes 2 each: 8 + 2(5) = 18 valence electrons

conclusions: this is a special case, where ionic counting is the same as neutral counting, all fragments being neutral. Since this is an 18-electron complex, it is expected to be isolable compound.

Ferrocene, (C5H5)2Fe, for the central Fe:

neutral counting: Fe contributes 8 electrons, the 2cyclopentadienyl-ringscontribute 5 each: 8 + 2(5) = 18 electrons

ionic counting: Fe2+contributes 6 electrons, the two aromatic cyclopentadienyl rings contribute 6 each: 6 + 2(6) = 18 valence electrons on iron.

conclusion: Ferrocene is expected to be an isolable compound.

These examples show the methods of electron counting, they are aformalism, and don't have anything to do withreal lifechemical transformations. Most of the 'fragments' mentioned above donotexist as such; they cannot be kept in a bottle: e.g. the neutral C, the tetraanionic C, the neutral Ti, and the tetracationic Ti are notfreespecies, they are always bound to something, for neutral C, it is commonly found in graphite, charcoal, diamond (sharing electrons with the neighboring carbons), as for Ti which can be found as its metal (where it shares its electrons with neighboring Ti atoms!), C4and Ti4+'exist' only with appropriate counterions (with which they probably share electrons). So these formalisms are only used to predict stabilities or properties of compounds!

Square Planar Complexes

In a square planar, there are four ligands as well. However, the difference is that the electrons of the ligands are only attracted to the \(xy\)plane. Any orbital in the xy plane has a higher energy level. There are four different energy levels for the square planar (from the highest energy level to the lowest energy level): dx2-y2, dxy, dz2, and both dxzand dyz.

INCLUDEPICTURE "http://chemwiki.ucdavis.edu/@api/deki/files/19106/2b.jpg?revision=1" \* MERGEFORMATINET Figure 6: Splitting of the degenerate d-orbitals (without a ligand field) due to an square planar ligand field.

The splitting energy (from highest orbital to lowest orbital) is \(\Delta_{sp}\) and tends to be larger then \(\Delta_{o}\)

\[\Delta_{sp} = 1.74\,\Delta_o \tag{2}\]

Moreover, \(\Delta_{sp}\) is also larger than the pairing energy, so the square planar complexes are usuallylow spincomplexes.

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