Key Concepts in Coordination Chemistry
The coordinate bond, also known as a dative bond, is a fundamental concept in coordination chemistry Unlike standard covalent bonds, which involve the sharing of electrons between two atoms, a coordinate bond forms when one atom donates a pair of electrons from a lone pair to an empty orbital of another atom This unique interaction highlights the distinct nature of molecular structures and their bonding mechanisms.
The concept of a covalent bond, introduced by G.N Lewis nearly a century ago, is fundamental to chemistry, providing insights into single, double, and triple bonds, as well as lone pairs of electrons This foundation led to the development of valence bond theory, an early quantum mechanical approach that articulated Lewis's ideas through wave-functions While these principles are integral to coordination chemistry, the field also requires the understanding of coordinate bond formation, where the bond pair of electrons comes solely from one atom In this context, the bonding interaction between an electron-pair acceptor (A) and an electron-pair donor (:D) can be represented as a simple equation.
The product can be represented as A←:D or A←D, indicating the direction of electron donation, or simply as A D when the bonding nature is clear This simplified representation is valid because, once formed, the covalent bond is indistinguishable from a standard covalent bond Additionally, the process is reversible; if the A D bond is broken, the lone pair of electrons originally donated by :D stays with that entity.
Coordination compounds feature a central atom or ion that forms coordinate bonds with multiple surrounding atoms, ions, or groups The central atom, typically a metal or metalloid, acts as an electron acceptor, while the surrounding species, known as ligands, donate lone pairs of electrons to empty orbitals on the central atom This interaction leads to the formation of a coordination compound, also referred to as a coordination complex or simply a complex Further exploration of these concepts will be provided below.
Introduction to Coordination Chemistry Geoffrey A Lawrance
C 2010 John Wiley & Sons, Ltd www.pdfgrip.com
A schematic view of ammonia acting as a donor ligand to a metalloid acceptor and to a metal ion acceptor to form coordinate bonds.
In coordination chemistry, the electron-pair donor, or ligand, is characterized by its ability to provide lone pairs of electrons to an electron-pair acceptor Ligands can range from simple monatomic ions, such as fluoride (F−), to complex polymers, but they all share the common feature of having one or more lone pairs on electronegative donor atoms Common donor atoms include heteroatoms like nitrogen (N), oxygen (O), and sulfur (S).
A wide variety of organic molecules, including halide ions and ammonia (NH3), can function as ligands due to their ability to donate electron pairs Ammonia, with its lone pair of electrons, serves as a classic example of a successful ligand Similarly, the carbanion (−CH3) is isoelectronic with ammonia and qualifies as a ligand under basic definitions Even hydrogen, in the form of its hydride (H−), possesses a pair of electrons that allows it to act as a ligand Ultimately, the essential factor is not the specific type of donor atom, but its capacity to supply an electron pair.
Coordinate covalent bonds are typically formed with acceptors that are metals or metalloids In the case of metalloids, the formation of these bonds generally leads to an increase in the number of groups or atoms bonded to the central atom, which can be explained through simple electron counting based on the donor–acceptor model However, with metal ions, this straightforward approach is less effective, as the number of new bonds formed during complexation does not always correspond to the apparent vacancies in the metal's valence shell, necessitating a more advanced model Notably, metal ions exhibit a strong tendency towards complexation, as 'naked' ions are rare, and complexation occurs even in the gaseous state This phenomenon illustrates that coordinate bond formation is energetically favorable, resulting in a system that is more stable than the individual components alone.
The anionic compound SiF6 2− exhibits a classical octahedral structure, commonly found in various metal complexes Positioned below carbon in the Periodic Table, silicon shares some limited chemical similarities with carbon Nonetheless, the simple valence bond theory and octet rule provide a fundamental framework for understanding these compounds.
The octahedral [SiF6] 2− molecular ion is formed through the overlap of p orbitals from six fluoride anions with empty hybrid orbitals on a Si(IV) cation, resulting in six equivalent covalent σ bonds Unlike carbon, traditional valence bond theory struggles to explain the formation of a silicon compound with six equivalent bonds The Si(IV) center, which has lost its four valence electrons, binds to six F− anions, each donating an electron pair to the silicon's empty orbitals Hybridization is essential for accommodating this structure, where silicon combines one 3s, three 3p, and two 3d orbitals to create six equivalent sp³d² hybrid orbitals, arranged to minimize repulsion and form an octahedral shape Each hybrid orbital forms a coordinate covalent bond with a fluoride ion, resulting in σ bonds characterized by electron density along the axis connecting the atomic centers The molecular shape is determined by the type and number of orbitals involved in hybridization, with variations leading to other geometries such as tetrahedral (sp³), linear (sp), and trigonal planar (sp²).
Coordinate bonds form when a central atom or ion with vacant orbitals is joined by ionic or neutral atoms or molecules that provide lone pairs of electrons, resulting in coordination compounds While the basic valence bonding model applies to metal ions, adjustments are necessary due to d orbital electrons, necessitating more advanced models Among these, molecular orbital theory stands out for its focus on the overlap of atomic orbitals of similar energy to create molecular orbitals for electron allocation Although this theory offers precise descriptions of molecular properties, its complexity and time-consuming nature make simpler models more commonly used, especially for larger complexes.
In Lewis' theory, the essential elements involve an empty orbital on one atom and a filled orbital with a lone pair of electrons on another Ligands that provide the lone pair are often classified as bases in the Brønsted–Lowry acid-base framework, as they can accept protons However, this discussion focuses on the interaction without proton transfer, emphasizing the acceptance of an electron-deficient species This leads to the broader definition of a Lewis base as an electron-pair donor and a Lewis acid as an electron-pair acceptor.
The H3B NH3 compound is typically classified as a coordination compound, formed by the interaction between the Lewis acid H3B and the Lewis base :NH3 Unlike the straightforward coordination of metal ions, the assembly of [H3C NH3]+ complicates the understanding due to traditional views on covalent bonding in carbon-based compounds This debate over the nature of the assembly limits the model's applicability for non-metals and metalloids, while coordination chemistry predominantly emphasizes metal compounds and their ions.
Coordination significantly impacts the new assembly, leading to structural changes characterized by variations in the number of bonds, bond angles, and distances This structural transformation is closely linked to alterations in the assembly's physical properties, which differ from those of its individual components In coordination complexes with metal atoms or ions at their core, even minor changes to the ligands can result in noticeable shifts in physical properties, such as color Advances in synthesis techniques and our understanding of physical methods allow for the tuning of these properties by varying ligands, enabling the achievement of specific outcomes like desired reduction potentials.
Coordination compounds exhibit a limited range of basic shapes, dictated by the central atom and its attached ligands The physical properties of these compounds are influenced by the characteristics of the central atom, the ligand set, and the overall molecular shape While many coordination compounds feature a single central atom, known as monomers, there is an expanding variety of compounds that contain two or more central atoms, which may be of the same or different types.
Central atoms are interconnected through direct atom-to-atom bonding or by ligands that connect to at least two central atoms simultaneously This bridging arrangement, where ligands link multiple central atoms, is more prevalent than direct bonding The resulting structures can often be viewed as a collection of monomer units, forming what is known as a polymer or, in cases with fewer linked units, an oligomer While this article primarily focuses on simple monomeric species, we will also introduce examples of larger linked compounds when relevant.
A Who’s Who of Metal Ions
Commoners and ‘Uncommoners’
Metals are often perceived as common due to their daily presence in various forms However, the term 'common' is relative; for instance, while iron is more abundant than gold in the Earth's crust, gold is still more prevalent than rhenium This illustrates that even well-known elements in the first row of the d-block of the Periodic Table vary in their abundance.
1 H hydrogen 3 Li lithium 11 Na sodium 19 K potassium 37 Rb rubidium 55 Cs caesium 87 Fr francium
21 Sc scandium 22 Ti titanium 23 V vanadium 24 Cr chromium 25 Mn manganese 26 Fe iron 27 Co cobalt 28 Ni nickel 29 Cu copper 30 Zn zinc 39 Y ytterbium 57 La lanthanum 89 Ac actinium
40 Zr zirconium 72 Hf hafnium 104 Rf rutherfordium
41 Nb niobium 74 W tungsten 105 Db dubnium
43 Tc technetium 75 Re rhenium 106 Sg seaborgium
44 Ru ruthenium 73 Ta tantalum 107 Bh bohrium
45 Rh rhodium 76 Os osmium 108 Hs hassium
46 Pd palladium 77 Ir iridium 109 Mt meitnerium
42 Mo molybdenum 78 Pt platinum 110 Ds darmstadtium
47 Ag silver 48 Cd cadmium 79 Au gold 111 Rg roentgenium
80 Hg mercury 112 Uub ununbium 93 Np neptunium
4 Be beryllium 12 Mg magnesium 20 Ca calcium 38 Sr strontium 56 Ba barium 88 Ra radium 58 Ce cerium 90 Th thorium
59 Pr praseodynium 60 Nd neodynium 61 Pm promethium 62 Sm samarium 63 Eu europium 64 Gd gadolinium 91 Pa protoactinium 92 U uranium
115 94 Pu plutonium 95 Am americium 96 Cm curium
65 Tb terbium 97 Bk berkelium 98 Cf californium
66 Dy dysprosium 67 Ho holmium 99 Es einsteinium
5 B boron 13 Al aluminium 31 Ga gallium
14 Si silicon 32 Ge germanium 33 As arsenic 49 In indium 81 Tl thallium 113
50 Sn tin 51 Sb antimony 82 Pb lead 83 Bi bismuth 84 Po polonium 117 116 114
7 N nitrogen 6 C carbon 8 O oxygen 36 Kr krypton 10 Ne neon
The abundance of metals in the Earth's crust varies widely, with iron being the most prevalent at 41,000 ppm, while other metals like cobalt are much rarer at 20 ppm Common metals such as copper and zinc, with abundances of 50 ppm and 75 ppm respectively, are not as abundant as they may seem The availability of an element is influenced by factors like its concentration in accessible ore bodies and its commercial value, rather than just its average crustal abundance Iron is extensively mined due to its rich ore deposits and significant commercial importance, while rhenium, the rarest transition metal, is primarily a by-product of mining other metals and has limited applications However, advancements in technology have made it possible to isolate and utilize all naturally occurring metals, as well as synthetic elements For instance, technetium, a synthetic d-block element used in medical imaging, and americium, a synthetic f-block element found in smoke detectors, demonstrate the commercial viability of both natural and man-made metals.
The insights gained from these observations significantly influence coordination chemistry, as every metallic element in the Periodic Table is now accessible for study Each element presents a distinct set of unique properties and behaviors, making them valuable for research and application in various fields.
The distinction among transition metals lies in their relative cost and availability, which has influenced a commercial focus on the coordination chemistry of lighter elements, ranging from vanadium to zinc Nature has similarly favored these lighter transition elements, as evidenced by their prevalent use in the active sites of most metalloenzymes.
Redefining Commoners
In addition to availability, oxidation states provide a significant chemical perspective on commonality among metals For the prevalent main group element aluminium, only the Al(III) oxidation state is significant In contrast, iron, the most common transition metal, exhibits both Fe(II) and Fe(III) as common oxidation states, with higher states like Fe(IV) being rare Conversely, rhenium shows a different pattern, as its high oxidation state, Re(VI), is common, while lower states such as Re(III) and Re(II) are infrequent These observations highlight that each metal may possess one or more typical oxidation states alongside a variety of less common ones, with some states being entirely inaccessible.
The uncommon oxidation states of metal ions are influenced by their surrounding environment, specifically the groups or atoms that coordinate with them The relationship between coordinating groups and the oxidation states a metal can exhibit is significant and will be examined further Additionally, the concept of 'common' oxidation states in metal complexes is dynamic within coordination chemistry, as it evolves based on ongoing research and discoveries; what was once considered 'unknown' may transition to 'very rare' over time.
‘uncommon’ as more chemists beaver away at extending the chemistry of an element At
Transition metal complexes exhibit various oxidation states, with d-electron counts corresponding to each state Common oxidation states are highlighted in black, while less common ones are shown in grey Figure 1.4 illustrates the oxidation states of first-row d-block elements, revealing that oxidation states two and three are the most prevalent Hydrated transition metal ions with charges exceeding 3+ are unstable in water, necessitating the presence of additional ligands Although definitions of common oxidation states may vary, the overall trends in oxidation states remain consistent.
Most metals exhibit a wide range of oxidation states, limited by the depletion of d electrons or by reaching a high reduction potential that destabilizes the ion This instability often results in immediate oxidation-reduction reactions that revert the metal to a more stable oxidation state As one moves from left to right across the Periodic Table, it becomes increasingly difficult to utilize all d electrons due to the increasing number of d electrons and reduced shielding from the nucleus's charge Nonetheless, coordination chemists still have ample options to explore.
The standard reduction potential (E 0) is a key indicator of a metal's stability in a specific oxidation state, measured against the standard hydrogen electrode (SHE) with a potential of 0.0 V A higher E 0 value signifies a greater difficulty for metal oxidation to occur, indicating that the metal is more stable in its higher oxidation state Conversely, a lower E 0 value suggests that the metal is more easily reduced to a lower oxidation state Additionally, metal activity is closely related to its reactivity with protic solvents, such as water, and correlates with electronegativity, particularly among very electropositive metals with significant cation reduction potentials.
Electropositive metals, characterized by low electronegativities (around -1.6 V), include the s block and all lanthanide metals, displaying cation reduction potentials up to approximately 0 V In contrast, electronegative metals, primarily found in the second and third rows of the d block, possess positive cation reduction potentials and are not corroded by oxygen, unlike their electropositive counterparts This variation in reactivity highlights the distinct behaviors of these metal classes in redox processes.
The definition of commonality among metal ions can be understood through the number of ligand donor groups attached to the central metal Historically, cobalt(III) was believed to consistently exhibit a coordination number of six, with six donor groups bonded to the metal ion While this remains the predominant coordination number for cobalt(III), stable examples of coordination numbers five and four have emerged In the case of cobalt(II), coordination numbers four and six are typical, leading to an increasing number of instances of the intermediate coordination number five Similarly, five-coordination in copper(II) has become nearly as common as four or six, highlighting the evolving nature of chemical definitions This reflects the broader challenge in chemistry, as the volume of research papers published annually continues to rise, making it increasingly difficult to keep up with the ever-changing landscape of the field.
Metals in Molecules
Metals in the Natural World
Metals in the Earth's crust are primarily found in inorganic environments, such as rocks and soils, both on land and underwater When geological processes lead to high concentrations of metals, these can form ore deposits, which are defined economically rather than scientifically Metals also exist as dissolved cations in water bodies, with notable concentrations like sodium in seawater Despite low concentrations, such as gold in seawater, the vast size of the oceans means significant amounts of metals are dispersed in aquatic environments Additionally, metals are found within living organisms, with transition metals like iron, zinc, and copper being predominant Some plants, like Hybanthus floribundus from Western Australia, can hyper-accumulate specific metals, such as nickel, in relatively high concentrations Overall, metal ion levels vary among species and environments, but they are generally present in trace amounts, as high levels can be toxic to living organisms, exemplified by ryegrass toxicity.
Cu⬎Ni⬎Mn⬎Pb⬎Cd⬎Zn⬎Al⬎Hg⬎Cr⬎Fe, with each species displaying a unique trend.
Metals are integral to various biomolecules in living organisms, where their presence is purposeful rather than coincidental They serve essential functions, ranging from creating an ionic environment to being critical components at active sites in enzymes The predominant metals found in biological systems are lighter alkali, alkaline earth, and transition elements.
Transition metals, particularly iron, copper, and zinc, are essential for the functioning of organisms, but nearly all first-row transition elements contribute to biological processes Additionally, even heavier elements like molybdenum and tungsten have been identified as having important roles in various biological functions.
Table 1.1 Typical concentrations (ppm) of selected metals ions in nature.
Metal Earth’s crust Oceans Plants (ryegrass) Animals (human blood)
Metal ions play crucial roles due to their high charge and surface charge density, which enhance their ability to form strong coordinate bonds with organic molecules Additionally, their capacity to vary oxidation states further contributes to their significance in biological systems A more detailed exploration of metals in biological environments will be provided in Chapter 8.
Metals in Contrived Environments
Over the past century, the field of chemistry has evolved significantly, particularly in our ability to design and synthesize millions of new molecules, marking the advent of the 'designer' molecule era Many of these organic molecules can bind to metal ions or be transformed into derivatives that can, leading to an essentially infinite variety of potential synthetic metal complexes The origins of synthetic metal complexes trace back to early examples like Prussian blue, a cyanide complex of iron, and hexaamminecobalt(III) chloride, discovered in 1798, which sparked interest in understanding their unique properties Initial identification of new compounds often relied on the names of their creators, resulting in colorful nomenclature such as Magnus’s green salt and Erdmann’s salt However, as the number of discovered compounds grew, a more systematic approach to naming emerged, moving away from color-based names to a structural nomenclature that remains in use today.
The fundamental principles of chemistry remain unchanged regardless of whether a metal complex is synthetic or natural, leading to new properties arising from novel assemblies rather than alterations in the underlying rules A prime example is Vitamin B12, a unique biomolecule centered on cobalt, known for its stability across three oxidation states: Co(III), Co(II), and Co(I), and featuring a C-Co(III) bond Initially, few low molecular weight synthetic cobalt(III) complexes exhibited stability in these oxidation states, suggesting that metals in biological systems were 'special.' However, advancements in synthetic chemistry have established stable cobalt complexes across all oxidation states and clarified the Co(III)-carbon bond with simple ligands The perceived uniqueness of metals in biology stems from their complex macromolecular environments, which, although challenging to replicate precisely in the lab, can be sufficiently mimicked to reflect certain chemical behaviors.
Synthetic coordination chemistry allows chemists to create compounds that surpass natural limitations, utilizing facilities and methods not found on Earth For instance, technetium, a radioactive element that has vanished from the planet, can now be produced in nuclear reactors, making its synthetic chemistry widely accessible Boron hydrides and carboranes, which are too reactive to exist geologically, showcase the rich chemistry derived from boron Additionally, nitrogen's ability to form carbon-based amines enhances metal ion binding, with synthetic amines surpassing natural counterparts Unlike natural processes, which occur under specific conditions, synthetic chemistry enables the design of novel molecules with tailored shapes and properties, leading to innovative applications and technologies limited only by human creativity.
Natural or Made-to-Measure Complexes
Metal complexes form naturally when metal ions interact with molecules that can bind to them, resulting in complexation When a metal salt is dissolved in water, the cations and anions become hydrated as they interact with water molecules, with the metal ion playing a crucial role in this process.
Lewis acids interact with water, acting as a Lewis base, leading to a defined coordination structure characterized by multiple M n + ←:OH2 bonds The experimentally determined M-O-H angle of approximately 130° supports the participation of a lone pair on the tetrahedral oxygen This coordinate bond is fundamental to both natural and synthetic complexes.
Metals are typically found in trace amounts within living organisms, yet advancements in isolation and concentration techniques have enabled the recovery of biological complexes The commercial availability of various metalloproteins highlights this capability, although sourcing some compounds from nature can be costly and restrictive To meet global demand, many drugs and natural compounds are now synthesized reliably and affordably This synthetic approach also extends to molecules that bind metal ions and their corresponding metal complexes Common over-the-counter metal compounds, such as zinc supplements, are often provided as synthetic zinc(II) amino acid complexes For further insights into biological coordination chemistry, refer to Chapter 8.
Hydrometallurgical processing, a water-based method, often utilizes complexation to isolate metal ions from ores, as seen in gold recovery where oxygen and cyanide ions form a soluble gold(I) cyanide complex Similarly, copper(II) ions are recovered through solvent extraction into kerosene and then back-extracted into aqueous acid for isolation via reduction In contrast, pyrometallurgical processes rely on high-temperature reduction reactions of oxide ores, while electrochemical processes, such as those used for aluminium and sodium recovery from molten salts, are also widely employed in industry.
The Road Ahead
In coordination chemistry, metals play a crucial role as central atoms, but it’s essential to recognize their partners, known as ligands A coordination complex can be likened to a molecular marriage, where both the metal and ligand contribute unique qualities, resulting in a union that is distinct from their individual components This analogy, while anthropomorphic, effectively illustrates that these complexes are not inert; they can undergo ligand exchange and alter their structure based on oxidation states and ligand preferences, reflecting a dynamic partnership The strength of this union can be experimentally measured, emphasizing the importance of compatibility between metals and ligands Ultimately, these concepts of partnership and compromise are fundamental to understanding coordination chemistry.
This book will explore ligands, metal-ligand assembly, and their implications, including molecular shape, stability, and properties, along with methods for measurement and interpretation We will also examine metal complexes in natural and commercial contexts, while considering future developments The goal is to provide a comprehensive overview suitable for an introductory text.
A coordination complex consists of a central atom, usually a metal ion, bound to a set of ligands by coordinate bonds.
A coordinate covalent bond is distinguished by the ligand donor atom donating both electrons (of a lone pair) to an empty orbital on the central atom to form the bond.
A ligand is a Lewis base (electron-pair donor), the central atom a Lewis acid (electron- pair acceptor).
A 'common' metal can be defined by its availability in the Earth's crust, but in coordination chemistry, it is more accurately described by factors such as its preferred oxidation state, the number of coordinated donors, and the types of donors it favors.
Metal ions may exist and form complexes in a number of oxidation states; this is particularly prevalent in the d block.
First row d-block metal ions are found dominantly in the M(II) or M(III) oxidation states Heavier members of the d block tend to prefer higher oxidation states.
For those seeking to strengthen their understanding of general chemistry, "Chemistry: Molecules, Matter and Change" by Atkins and Jones (2000) serves as an excellent introductory textbook that reviews fundamental concepts Additionally, "The Periodic Table at a Glance" by Beckett and Platt (2006) offers a concise, well-illustrated overview of periodicity in inorganic chemistry, making it an ideal resource for undergraduate students.
Gillespie, R.J and Popelier, P.L.A (2002) Chemical Bonding and Molecular Geometry, Oxford University Press A coverage from the fundamental level upward of various models of molecular bonding.
Housecroft and Sharpe's "Inorganic Chemistry" (3rd ed., 2008) is a well-crafted and visually appealing advanced textbook that serves as a valuable resource for students and professionals in the field of inorganic chemistry.
Membership: Being a Ligand
What Makes a Ligand?
Ligands encompass a wide variety of molecules, including inorganic atoms, ions, and organic compounds, capable of binding to metal ions While metals can form metal–metal bonds, we will not consider them as ligands in this context The number of molecules that can participate in ligation is vast, with most organic molecules either acting as ligands directly or being convertible into ligands The fundamental requirement for a molecule to function as a ligand is the presence of at least one lone pair of electrons, with the atom carrying the lone pair referred to as the donor atom When this atom is part of a functional group, such as an amine or carboxylate, it is called a donor group, although typically only one specific donor atom from the group binds to the metal Initially, ligands were viewed as passive participants, primarily serving to donate electron pairs rather than actively engaging in the bonding process.
We now know that ligands influence the central metal ion significantly and can display reactivity and undergo chemistry of their own while still bound to the metal ion, which
Introduction to Coordination Chemistry Geoffrey A Lawrance
C 2010 John Wiley & Sons, Ltd www.pdfgrip.com
Metal ions are hardly ever found naked They are always clothed with ligands.
An anthropomorphic view of being a metal coordination complex. we shall return to in Chapter 6; this does not alter the fundamentals of their behaviour as ligands, however.
Ligands (often represented by the general symbol L) may present a single donor atom with a lone pair for binding to a metal ion and thus occupy only one coordination site,
A monodentate ligand is defined as a ligand that binds to a metal ion through a single donor atom, with classical examples including ammonia (NH3), water (H2O), and chloride ion (Cl−) While water and chloride possess multiple lone pairs, they primarily act as monodentate ligands Many ligands, however, can function as polydentate ligands, offering multiple donor groups capable of binding to the same metal ion Typically, heteroatoms such as oxygen, nitrogen, sulfur, and phosphorus in organic molecules possess lone pairs of electrons, making them suitable donor atoms in ligands Identifying these heteroatoms is crucial for determining a molecule's potential as a ligand and the number of donor groups available for binding to a metal ion Additionally, the spatial arrangement and local environment of these donor atoms within a larger molecule influence their binding capabilities.
‘available’ donor groups may bind collectively to a single metal ion, but identification of their presence is always the first step.
Making Attachments – Coordination
In the gas phase, a ligand can encounter a 'naked' metal or ion, but this scenario is rare In liquid or solid states, ligands typically interact with metals that are already surrounded by other ligands, often including solvent molecules that can function as ligands as well The binding of a new ligand to the inner coordination sphere of a metal ion generally necessitates the replacement of an existing ligand, a process known as substitution, which will be further explored in Chapter 5.
The complexation of ‘naked’ metal ions, such as copper(II), is evident when a hydrated salt like copper(II) sulfate (CuSO4·5H2O) is dried to form an anhydrous salt (CuSO4) Upon dissolving this anhydrous salt in water, a pale blue solution is formed as water molecules rapidly bind to the metal ion as ligands, a process associated with the heat of hydration When the solvent is evaporated, the hydrated salt CuSO4·5H2O is recovered, indicating that water molecules are tightly bound to the copper ion This can be demonstrated by measuring weight changes during heating, revealing that complete removal of water requires temperatures over 200 °C Advanced experimental methods allow for the observation of ions in solution, confirming that the copper ion is surrounded by a well-defined sheath of water molecules, forming an inner coordination sphere through coordinate covalent bonds This inner sphere is further surrounded by an outer coordination sphere of hydrogen-bonded water molecules, with additional layers forming until they blend into the bulk water, resulting in progressively diminishing electrostatic influence from the metal ion.
The lifetime of a complex ion is influenced by the dynamics of ligand exchange, which involves the swapping of water molecules within the coordination sphere of an aquated metal ion in pure water Although one might assume that a complex ion retains the same set of water molecules indefinitely, it is common for outer coordination sphere water molecules to replace those in the inner sphere This process is subtle and often undetectable, akin to replacing a cold can with a warm one without anyone noticing By introducing water with a different oxygen isotope, researchers can study ligand exchange at the molecular level The rate of this exchange varies significantly depending on the type and oxidation state of the metal ion, as well as the ligands involved, leading to complex ions with either extreme longevity or a relatively fixed coordination sphere The concept of ligand exchange will be further explored in Chapter 5.
Putting the Bite on Metals – Chelation
Ammonia is a classic simple ligand due to its single lone pair of electrons, which limits it to forming only one coordinate covalent bond In contrast, a water molecule contains two lone pairs on its oxygen atom but typically forms just one coordinate bond This limitation arises from the spatial arrangement of the lone pairs; after one bond is formed, the remaining lone pair is oriented in a way that prevents it from coordinating with the same metal ion Instead, it can only achieve coordination through a different metal, a scenario referred to as bridging We will further explore the potential for this bridging phenomenon with water molecules later.
O H M' H free coordinated free coordinated bridging
Free and coordinated ammonia and water molecules exhibit distinct orientations; the second lone pair of water is positioned to prevent interaction with the same metal center as the first However, this lone pair could theoretically bind to a second metal center, facilitating a bridging coordination mode.
To facilitate the interaction of two lone pairs with a single metal ion, we can position them on different atoms, using two ammonia residues connected by a carbon atom as an example Although this configuration is not highly stable, it serves our purpose Initially, either lone pair from the nitrogen atoms can bond with the metal ion The second amine group can rotate around the fixed M-N-C structure, allowing its lone pair to align more favorably towards the metal ion If the covalent bonds are slightly distorted, it is possible for both lone pairs to coordinate with the same metal ion effectively.
The carboxylate group (R COO−) serves as a stable example of coordination chemistry, capable of binding to metals in three distinct ways: through a single oxygen atom to one metal, bridging two metals with each attached to one oxygen, or coordinating to one metal via both oxygen atoms While the size of the atom ring, which includes the metal and donor atoms, remains consistent across different coordination modes, the types of donor atoms vary.
Chelation occurs when a single ligand uses two different donor atoms to bond with the same metal ion, resulting in the formation of a chelate ring This term, introduced by Morgan and Drew in 1920, likens the process to a lobster using both claws to grasp its prey, highlighting that chelates typically create stronger complexes compared to simple monodentate ligands A chelate ring is defined as a cyclic structure comprising the two donor atoms, the metal ion, and the connecting portion of the ligand framework The size of the chelate ring is determined by counting the number of covalently linked atoms.
Diaminomethane is a molecule featuring two amine groups, allowing for unique bonding capabilities When the first amine group coordinates, the second amine's lone pair can align more effectively for bonding, minimizing bond angle distortion This results in a chelated coordination mode, enhancing the molecule's stability and reactivity.
- - free monodentate didentate bridging didendate chelate
Metal ion binding to carboxylate groups can occur through various coordination modes, including monodentate and bidentate interactions In a continuous ring structure that begins and ends at the metal ion, the chelation process is exemplified by the sequence M→O→C→O→, as illustrated in Figure 2.4.
O returning us to the M, so four atoms are involved in the continuous ring, meaning it is a four-membered chelate ring.
Diaminomethane, depicted in Figure 2.3, forms a four-membered chelate ring similar to the previously discussed carboxylate This ring consists of four atoms, including a metal, connected by four covalent bonds, two of which are coordinate bonds Just as organic compounds with specific ring sizes exhibit inherent stability, chelation significantly enhances the stability of metal complexes with chelate rings of certain dimensions.
Using diaminoethane, also known as ethane-1,2-diamine or ethylenediamine, instead of diaminomethane results in the formation of a five-membered chelate ring through chelation This process begins with one nitrogen atom bonding to the metal, followed by the rotation of the remaining lone pair to achieve the correct orientation for effective binding The initial bonding of the first nitrogen to the metal ensures that the second nitrogen remains in close proximity, which facilitates its coordination and enhances the chelation process.
In diaminoethane, the two amines must be in a cis configuration for effective chelation, while a trans arrangement leads to bridging with two separate metals The flexible nature of this ligand allows for easy rotation around the C-C bond, enabling a seamless transition between different conformations in the unbound ligand.
The stepwise process for chelation of diaminoethane This features initial monodentate formation,rearrangement and orientation of the second lone pair, and its subsequent binding to form the chelate ring.
Diaminoethane exhibits the freedom to rotate around the carbon-carbon bond, allowing for the formation of cis and trans isomers, which can engage in chelation and bridging, respectively In contrast, the rigid structure of diaminobenzene prevents rearrangement, resulting in two distinct isomers that serve unique coordinating functions as didentate ligands.
Rigid ligands like diaminobenzene exhibit distinct structural differences between their isomers; the 1,4- (para or trans) isomer serves only as a bridging ligand, while the 1,2- (ortho or cis) isomer can chelate, both classified as didentate ligands due to their two nitrogen donor sites The chelate ring formed with 1,2-diaminobenzene remains flat due to the influence of its rigid aromatic structure, in contrast to the non-planar ring of diaminoethane, where the nitrogen and carbon centers maintain their tetrahedral geometry This results in a puckered ring shape, as illustrated by the orientation of the nitrogen-methylene-nitrogen plane Additionally, enforcing planarity through sp²-hybridized carbon, as seen in carboxylates, can influence the overall structure of the ligand.
Chelate ring conformations in chelated diaminoethane (ethane-1,2-diamine, en), designated as␦and
The article discusses the structural views of a chelate ring involving the glycine anion (H2NCH2COO−) and the oxalate dianion (−OOC COO−) It highlights that the glycine anion forms a five-membered chelate ring characterized by a tetrahedral and a trigonal planar carbon, exhibiting less puckering compared to diaminoethane In contrast, the oxalate dianion, with two trigonal planar carbons, maintains a completely flat configuration in its chelated form Additionally, the article notes that the two conformations of the glycine anion are mirror images of each other, allowing for easy interconversion due to the minimal energy barrier between them.
The puckered diaminoethane exhibits distinct conformations due to its rigid chelate ring, which causes the protons on each carbon to be nonequivalent; one proton is oriented axially (Hax) while the other is positioned equatorially (Heq) Despite this rigidity, the molecule retains enough flexibility to undergo inversion, allowing one carbon to move upwards and the other downwards, resulting in two mirror-image forms known as and ␦ This phenomenon of having non-flat chelate rings leads to the existence of such conformers.
Didentate chelates come in a wide variety, making them complex to navigate However, many classical examples form flat chelate rings due to the shape of the donor groups or the enforced planarity from conjugation Figure 2.8 illustrates common ligands along with their abbreviated names Notably, the chain of atoms connecting the donor atoms can differ, resulting in various chelate ring sizes; nevertheless, four-atom chains leading to five-membered chelate rings are the most prevalent This topic will be further explored in the next section.
- O oxalato ox 2- glycinato ethane-1,2-diamine
2,4-dioxopentane-3-ido (or acetylacetonato) acac
Some common didentate chelating ligands Common abbreviations used for the ligands are given to the right of each line drawing.
L L bond length change bond angle change adjustment
Do I Look Big on That? – Chelate Ring Size
As the length of the atomic chain connecting donor atoms increases, the size of the chelate ring also expands, influencing the 'bite' of the chelate and its stability Adjustments in bond distances and angles within the organic framework of a ligand require more energy compared to those around the metal center, leading to more significant changes near the metal When mismatches in chelation arise, variations in metal-ligand bond lengths and ligand-metal-ligand angles are primarily used to accommodate the nonideal fit, although some minor adjustments may still occur within the organic ligand.
There is no definitive upper limit on chelate ring size, but excessively long chains between donor atoms can reduce the stability and effectiveness of the complex, making chelation less efficient As the distance between the second donor and the anchored first donor increases, the likelihood of successful binding diminishes Figure 2.10 illustrates examples of chelate rings ranging from three to seven members, highlighting how the experimentally measured L M L angle varies with ring size While simple bidentate ligands are sometimes classified by the number of atoms in the chain separating the donor groups—such as 1,1-ligands for four-membered rings, 1,2-ligands for five-membered rings, and 1,3-ligands for six-membered rings—this classification is not commonly used in practice.
Research indicates that there is an optimal chelate ring size that enhances the stability of metal-ligand complexes Initially, as the ring size increases, the stability of the assembled complex improves; however, this stability eventually decreases with further increases in ring size This trend is influenced by several factors, including the type of metal ion, the donor groups present, and the ligand framework utilized Notably, this pattern is consistently observed among lighter metals in the first row of the periodic d-block.
3-membered⬍4-membered⬍5-membered⬎6-membered⬎7-membered.
The preference for five-membered chelate rings is evident, as demonstrated by experimental measurements of O-donor series, illustrated in Figure 2.10 This trend is reflected in the stability of metal complexes, highlighting the significance of chelate ring size in coordination chemistry.
3 4 5 6 7 dioxygen carbonate oxalate malonate succinate
Chelate rings, ranging from three to seven members, significantly influence the L M L intraligand angle, as illustrated in the accompanying figure This relationship highlights the tendency of ligands, such as oxalate, malonate, and succinate, to form complexes with various divalent transition metal ions, demonstrating the importance of ring size in coordination chemistry.
The trend in ring sizes is consistently observed as 5-membered rings being preferred over 6-membered and 7-membered rings While this trend holds for lighter metal ions, larger metal ions may favor different ring sizes due to their size and preferred bond lengths Additionally, there is a notable variation with different metals, regardless of the ligand involved, which will be discussed further.
Different Tribes – Donor Group Variation
Different types of donors exist, as illustrated by various molecules in which nitrogen (N), oxygen (O), sulfur (S), phosphorus (P), and even carbon (C) atoms form bonds.
Mn Fe Co Ni Cu Zn mal (6) log K 1 suc (7) ox (5)
The stability of metal(II) ion complexes is influenced by the size of the chelate ring formed by O,O-chelates such as oxalate (5-membered), malonate (6-membered), and succinate (7-membered) Just as different human tribes exhibit unique characteristics, ligands with varying donor atoms also display distinct properties The donor atom, being directly bonded to the central metal, significantly impacts the metal's behavior, affecting the overall chemical and physical properties of the complex.
Most donor atoms in ligands belong to the main group (p block) of the Periodic Table, with several common simple ligands frequently encountered in various applications.
Ligands featuring donor atoms such as nitrogen (N), oxygen (O), phosphorus (P), sulfur (S), and halogen anions are prevalent in various chemical contexts These donor atoms represent the majority of ligands encountered, including those found in natural biomolecules Additionally, other elements like carbon also serve as ligands, expanding the diversity of coordination chemistry.
H3C − , for example, is an effective donor, and there is an area of coordination chemistry (organometallic chemistry) devoted to compounds that include M C bonds, addressed later in Section 2.5.
Ligands frequently possess multiple donor groups, leading to the formation of mixed-donor ligands, which are prevalent in coordination chemistry A prime example of this is amino acids, such as [H2N CH(R) COOH], which feature both nitrogen (amine) and oxygen (carboxylate) donor sites Given the variety of donor groups available to a metal ion, it is expected that certain preferences may arise, a topic that will be explored further in Chapter 3.5.
Ligands with More Bite – Denticity
Ligands can attach to metals at one or multiple coordination sites, with some capable of fully satisfying the coordination sphere of a metal ion or even binding several metal ions The concept of denticity defines the number of coordinated donor groups in a ligand, indicating how many donor groups are bound to a metal ion Monodentate ligands have one donor group, didentate ligands have two, tridentate ligands have three, and tetradentate ligands have four This terminology simplifies discussions about ligand binding, allowing for concise references like "a didentate ligand" instead of describing the specific donor groups involved.
Denticity and the number of potential donor groups in a ligand are not synonymous For instance, the amino acid glycine can interact with a metal in multiple ways: through its carboxylate group, its amine group, or as a chelate by utilizing both groups simultaneously The first two binding modes illustrate the concept of denticity effectively.
Possible coordination modes for the amino acid ligand glycine. the molecule acting as a monodentate ligand, the last, of the molecule acting as a didentate chelate ligand (Figure 2.12).
The binding of ligands to metals is influenced by the shape and positioning of the ligand within its framework Figure 2.13 illustrates various ligands with different denticities, focusing on a simple flat square shape for clarity It is important to recognize that other geometric configurations can also accommodate ligands, and some ligands can act as monodentate, highlighting the versatility of ligand coordination.
Ligands exhibit varying denticity, with many requiring deprotonation to form an anionic state for effective coordination, as this enhances their ability to donate electrons to cationic metals The capacity for deprotonation is influenced by the acidity of the donor groups, and not all groups can easily achieve this Additionally, some ligands possess multiple heteroatoms that can serve as donor sites, such as carboxylate (COO−) and amido (N−CO) groups However, the coordination of both donor sites to a single metal ion may be restricted by the ligand's molecular shape or may only occur under specific synthetic conditions Further exploration of polydentate ligands will be provided in Section 2.3.