Hubbry Logo
Coordinate covalent bondCoordinate covalent bondMain
Open search
Coordinate covalent bond
Community hub
Coordinate covalent bond
logo
7 pages, 0 posts
0 subscribers
Be the first to start a discussion here.
Be the first to start a discussion here.
Coordinate covalent bond
Coordinate covalent bond
Coordinate covalent bond
View on Wikipedia
from Wikipedia

In coordination chemistry, a coordinate covalent bond,[1] also known as a dative bond,[2] dipolar bond,[1] or coordinate bond[3] is a kind of two-center, two-electron covalent bond in which the two electrons derive from the same atom. The bonding of metal ions to ligands involves this kind of interaction.[4] This type of interaction is central to Lewis acid–base theory.

Coordinate bonds are commonly found in coordination compounds.[5]

Examples

[edit]
Formation of an adduct of ammonia and boron trifluoride, involving formation of a coordinate covalent bond.

Coordinate covalent bonding is ubiquitous.[6] In all metal aquo-complexes [M(H2O)n]m+, the bonding between water and the metal cation is described as a coordinate covalent bond. Metal-ligand interactions in most organometallic compounds and most coordination compounds are described similarly.

The term dipolar bond is used in organic chemistry for compounds such as amine oxides for which the electronic structure can be described in terms of the basic amine donating two electrons to an oxygen atom.

R
3N → O

The arrow → indicates that both electrons in the bond originate from the amine moiety. In a standard covalent bond each atom contributes one electron. Therefore, an alternative description is that the amine gives away one electron to the oxygen atom, which is then used, with the remaining unpaired electron on the nitrogen atom, to form a standard covalent bond. The process of transferring the electron from nitrogen to oxygen creates formal charges, so the electronic structure may also be depicted as

R
3N+
O−
Hexamminecobalt(III) chloride

This electronic structure has an electric dipole, hence the name polar bond. In reality, the atoms carry partial charges; the more electronegative atom of the two involved in the bond will usually carry a partial negative charge. One exception to this is carbon monoxide. In this case, the carbon atom carries the partial negative charge although it is less electronegative than oxygen.

An example of a dative covalent bond is provided by the interaction between a molecule of ammonia, a Lewis base with a lone pair of electrons on the nitrogen atom, and boron trifluoride, a Lewis acid by virtue of the boron atom having an incomplete octet of electrons. In forming the adduct, the boron atom attains an octet configuration.

The electronic structure of a coordination complex can be described in terms of the set of ligands each donating a pair of electrons to a metal centre. For example, in hexamminecobalt(III) chloride, each ammonia ligand donates its lone pair of electrons to the cobalt(III) ion. In this case, the bonds formed are described as coordinate bonds. In the Covalent Bond Classification (CBC) method, ligands that form coordinate covalent bonds with a central atom are classed as L-type, while those that form normal covalent bonds are classed as X-type.

Comparison with other electron-sharing modes

[edit]

In all cases, the bond, whether dative or "normal" electron-sharing, is a covalent bond. In common usage, the prefix dipolar, dative or coordinate merely serves to indicate the origin of the electrons used in creating the bond. For example, F3B ← O(C2H5)2 ("boron trifluoride (diethyl) etherate") is prepared from BF3 and :O(C2H5)2, as opposed to the radical species [•BF3]– and [•O(C2H5)2]+. The dative bond is also a convenience in terms of notation, as formal charges are avoided: we can write D: + []A ⇌ D → A rather than D+–A– (here : and [] represent the lone-pair and empty orbital on the electron-pair donor D and acceptor A, respectively). The notation is sometimes used even when the Lewis acid-base reaction involved is only notional (e.g., the sulfoxide R2S → O is rarely if ever made by reacting the sulfide R2S with atomic oxygen O). Thus, most chemists do not make any claim with respect to the properties of the bond when choosing one notation over the other (formal charges vs. arrow bond).

It is generally true, however, that bonds depicted this way are polar covalent, sometimes strongly so, and some authors claim that there are genuine differences in the properties of a dative bond and electron-sharing bond and suggest that showing a dative bond is more appropriate in particular situations. As far back as 1989, Haaland characterized dative bonds as bonds that are (i) weak and long; (ii) with only a small degree of charge-transfer taking place during bond formation; and (iii) whose preferred mode of dissociation in the gas phase (or low ε inert solvent) is heterolytic rather than homolytic.[7] The ammonia-borane adduct (H3N → BH3) is given as a classic example: the bond is weak, with a dissociation energy of 31 kcal/mol (cf. 90 kcal/mol for ethane), and long, at 166 pm (cf. 153 pm for ethane), and the molecule possesses a dipole moment of 5.2 D that implies a transfer of only 0.2 e– from nitrogen to boron. The heterolytic dissociation of H3N → BH3 is estimated to require 27 kcal/mol, confirming that heterolysis into ammonia and borane is more favorable than homolysis into radical cation and radical anion. However, aside from clear-cut examples, there is considerable dispute as to when a particular compound qualifies and, thus, the overall prevalence of dative bonding (with respect to an author's preferred definition). Computational chemists have suggested quantitative criteria to distinguish between the two "types" of bonding.[8][9][10]

Some non-obvious examples where dative bonding is claimed to be important include carbon suboxide (O≡C → C0 ← C≡O), tetraaminoallenes (described using dative bond language as "carbodicarbenes"; (R2N)2C → C0 ← C(NR2)2), the Ramirez carbodiphosphorane (Ph3P → C0 ← PPh3), and bis(triphenylphosphine)iminium cation (Ph3P → N+ ← PPh3), all of which exhibit considerably bent equilibrium geometries, though with a shallow barrier to bending. Simple application of the normal rules for drawing Lewis structures by maximizing bonding (using electron-sharing bonds) and minimizing formal charges would predict heterocumulene structures, and therefore linear geometries, for each of these compounds. Thus, these molecules are claimed to be better modeled as coordination complexes of :C: (carbon(0) or carbone) or :N:+ (mononitrogen cation) with CO, PPh3, or N-heterocycliccarbenes as ligands, the lone-pairs on the central atom accounting for the bent geometry. However, the usefulness of this view is disputed.[9][10]

References

[edit]
Revisions and contributorsEdit on WikipediaRead on Wikipedia

Coordinate covalent bond

View on Grokipedia
from Grokipedia
A coordinate covalent bond, also known as a dative covalent bond or coordination bond, is a type of covalent chemical bond in which the shared pair of electrons is supplied entirely by one of the two atoms, typically a Lewis base donating a lone pair to a Lewis acid. This bond forms through the interaction between an electron-rich atom (the donor) and an electron-deficient atom (the acceptor), resulting in a stable molecular entity where the bond's characteristics become indistinguishable from those of a regular covalent bond once established. Unlike conventional covalent bonds, where each atom contributes one electron to the pair, the coordinate bond originates from a single atom's lone pair, often represented by an arrow pointing from the donor to the acceptor in Lewis structures. Coordinate covalent bonds are fundamental in coordination chemistry, where they link ligands—molecules or ions with available lone pairs—to a central metal atom or ion, forming coordination complexes essential for applications in catalysis, bioinorganic chemistry, and materials science. Common examples include the ammonium ion (NH₄⁺), formed when ammonia (NH₃) donates its lone pair to a proton (H⁺), and the hydronium ion (H₃O⁺), resulting from water's donation to H⁺. In metal complexes like [Co(NH₃)₆]³⁺, multiple coordinate bonds from ammonia ligands surround the cobalt ion, stabilizing the structure and influencing its reactivity. These bonds also appear in simple molecules such as nitric acid (HNO₃), where the hydroxyl oxygen donates to the nitrogen, and carbon monoxide (CO), with the carbon-oxygen bond exhibiting partial dative character. The concept of coordinate covalent bonding emerged from Lewis acid-base theory, providing a framework to explain electron-deficient compounds and octet rule exceptions, such as boron trifluoride (BF₃) adducts. The term "dative bond" is a synonym for "coordination bond," highlighting the directional electron donation, and such bonds often exhibit greater polarity, greater bond lengths, lesser strength, and susceptibility to heterolytic cleavage compared to symmetric covalent bonds. In coordination compounds, the number of these bonds defines the coordination number, which dictates the geometry—such as octahedral or tetrahedral—and properties like color and magnetism of the complex.

Fundamentals

Definition

A coordinate covalent bond, also known as a dative bond, is a type of covalent bond in which both electrons in the shared pair are supplied by a single atom, referred to as the donor or Lewis base, to form a bond with an electron-deficient atom called the acceptor or Lewis acid. This distinguishes it from typical covalent bonds, where each atom contributes one electron to the pair. The dative of this bond emphasizes the unidirectional of the , with the acceptor providing no electrons initially but forming a bond through the . Once formed, the bond behaves similarly to other covalent bonds in terms of electron delocalization and strength. In Lewis structures, coordinate covalent bonds are conventionally denoted by an from the donor atom to the acceptor, such as A → B, to illustrate the direction of . This notation highlights the bond's role as a prerequisite for understanding mechanisms in chemical bonding, foundational to Lewis acid-base theory.

Historical Context

The concept of the coordinate covalent bond traces its origins to the late 19th and early 20th centuries, building on efforts to explain the structures of coordination compounds. In 1893, Alfred Werner proposed the idea of "coordinate valence" to account for the bonding in metal complexes, distinguishing between primary (ionizable) and secondary (non-ionizable) valences, where the latter represented attachments to a central metal atom without altering its charge. This framework laid the groundwork for understanding how ligands bind to metals, influencing the field of inorganic chemistry profoundly. Werner's theories were validated through his studies on optical isomers of coordination compounds, earning him the Nobel Prize in Chemistry in 1913 for his work on atomic linkages in molecules. The modern electron-pair basis for coordinate bonds emerged with Gilbert N. Lewis's seminal 1916 paper, "The Atom and the Molecule," where he introduced the octet rule and described covalent bonds as shared electron pairs. Lewis specifically differentiated coordinate (or semipolar) bonds, in which one atom donates both electrons of the pair while the other provides an empty orbital, as a subset of covalent bonding essential for explaining structures like ammonium ions. This electron-sharing model revolutionized valence theory, shifting focus from classical electrostatics to quantum-compatible descriptions. In the 1920s, Nevil V. Sidgwick refined and popularized these ideas, particularly in coordination chemistry, through his 1927 book The Electronic Theory of Valency. Sidgwick introduced the term "coordinate link" to describe dative bonds where the electron pair originates from one atom, applying it extensively to metal-ligand interactions and emphasizing its role in achieving stable electron configurations. His work bridged Lewis's foundational concepts with practical applications in complex compounds, solidifying the theoretical underpinnings. Over time, terminology evolved from Werner's "coordinate valence" and Sidgwick's "coordinate link" to the contemporary "coordinate covalent bond" or "dative bond," reflecting a unified view of these as covalent interactions with directional electron donation. The IUPAC now considers "coordinate covalence" and "coordinate link" obsolescent synonyms, while "dative bond" is deemed obsolete, standardizing usage in modern chemical nomenclature. This terminological progression underscored the bond's centrality in coordination chemistry, enabling advances in understanding molecular geometries and reactivity.

Formation Mechanism

Lewis Acid-Base Theory

In the Lewis acid-base theory, formulated by Gilbert N. Lewis in 1923, a Lewis acid is defined as a species that can accept an electron pair, typically possessing an empty orbital or deficient electron shell, while a Lewis base is a species that can donate an electron pair, often featuring a lone pair of electrons. This electron-pair perspective expands beyond earlier definitions, emphasizing the role of electron sharing in acid-base interactions rather than solely ionic or proton-based mechanisms. The formation of a coordinate covalent bond, also known as a dative bond, occurs when a Lewis base donates its lone pair to the empty orbital of a Lewis acid, resulting in a shared electron pair that constitutes the bond. In this process, both electrons originate from the base, distinguishing it from typical covalent bonds where electrons are contributed equally by each atom, yet the bond shares similar strength and characteristics once formed. This mechanism underlies the creation of adducts, stable molecular entities resulting from the acid-base union. Common examples of Lewis acids include (BF₃), which has an incomplete octet and an empty p-orbital, and the hydronium ion (H⁺), a simple electron-pair acceptor. Lewis bases such as (NH₃) and (H₂O) possess available lone pairs on or oxygen atoms, respectively, donation to acids. The general reaction for coordinate bond formation is represented as: Lewis base+Lewis acid→adduct\text{Lewis base} + \text{Lewis acid} \rightarrow \text{adduct} A classic illustration is: NH3+BF3→H3N→BF3\mathrm{NH_3 + BF_3 \rightarrow H_3N \rightarrow BF_3} where the arrow indicates the dative nature of the bond from nitrogen to boron. This theory differs from the Brønsted-Lowry framework, which limits acids to proton (H⁺) donors and bases to proton acceptors, often involving proton transfer in aqueous solutions. In contrast, Lewis theory accommodates coordinate bond formation without any proton involvement, applying to non-protonic solvents and a wider array of molecular interactions, thus providing a more comprehensive model for understanding dative bonding in coordination chemistry.

Orbital Interactions

In coordinate covalent bonds, the formation arises from the overlap between a filled orbital on the electron donor and an empty orbital on the electron acceptor. For instance, in the BF₃-NH₃ adduct, the empty p-orbital on the boron atom of BF₃ accepts the lone pair from the nitrogen atom of NH₃, resulting in a sigma-type interaction that shares the electron pair between the two centers. This overlap is facilitated by the directional properties of atomic orbitals, ensuring effective electron density accumulation along the internuclear axis. Hybridization plays a crucial role in optimizing this orbital alignment. In the donor, such as NH₃, the nitrogen atom adopts sp³ hybridization, positioning the lone pair in a hybrid orbital oriented for maximum overlap with the acceptor's empty orbital. On the acceptor side, BF₃ features boron in sp² hybridization, with the three sp² orbitals forming bonds to fluorine atoms and leaving the perpendicular p-orbital vacant for donation acceptance; upon complexation, the boron rehybridizes toward sp³ to accommodate the new bond. This adjustment enhances the geometric compatibility and strengthens the directional overlap. From a perspective, the donor-acceptor interaction combines the donor's highest occupied (HOMO), typically a lone-pair orbital, with the acceptor's lowest unoccupied molecular orbital (LUMO), an empty orbital, to form new and antibonding molecular orbitals delocalized over both fragments. In the BF₃-NH₃ , the primary donation occurs from the NH₃ HOMO (2a₁ symmetry) to the BF₃ LUMO (5a₁* symmetry), with a smaller back-donation from BF₃ filled orbitals to NH₃ empty orbitals, leading to a net charge transfer of approximately 0.26 electrons and stabilizing the complex through partial filling of the bonding orbital. This results in a sigma bonding orbital lower in energy than the isolated atomic orbitals, while the antibonding counterpart remains higher. The bond order in coordinate covalent bonds is typically 1, reflecting a single sigma bond from the two-center, two-electron interaction. However, in transition metal complexes, additional pi-backbonding from metal d-orbitals to ligand pi* orbitals can contribute to higher bond multiplicity, effectively increasing the bond order beyond 1 through synergistic sigma donation and pi acceptance. Conceptually, the orbital overlap can be visualized as the symmetric approach of the donor's lone-pair orbital (often s-p hybrid) toward the acceptor's empty p- or hybrid orbital, forming a sigma bond with cylindrical symmetry around the bond axis; this is analogous to a standard covalent sigma bond but initiated unilaterally from the donor. The resulting molecular orbital diagram shows the bonding combination below the non-bonding levels, emphasizing the stabilizing electron delocalization.

Properties

Bond Strength and Stability

The strength of coordinate covalent bonds is quantified by their bond dissociation energies (BDEs), which typically range from 50 to 200 kJ/mol, rendering them weaker than standard covalent bonds such as the C-N single bond at approximately 305 kJ/mol. This reduced strength arises from the directional donation of electron density from the Lewis base to the Lewis acid, resulting in a more polar character compared to symmetric covalent bonds. For instance, the BDE for the N-B bond in the ammonia-borane adduct (H₃N→BH₃) is about 114 kJ/mol, while that in the ammonia-boron trifluoride adduct (H₃N→BF₃) is lower at 92 kJ/mol, illustrating variability across adducts. Several factors influence the strength of coordinate bonds. The difference in electronegativity between the donor and acceptor atoms modulates the polarity and thus the electrostatic contribution to bond stability, with greater differences often leading to weaker bonds due to reduced orbital overlap efficiency. Orbital overlap efficiency, determined by the size, shape, and energy matching of the donor lone-pair orbital and acceptor empty orbital, is critical; optimal overlap, as in cases with similar orbital energies, enhances bond strength. In transition metal complexes, π-back-donation from metal d-orbitals to ligand π* orbitals further strengthens the coordinate bond by providing additional electron density delocalization. Coordinate bonds in adducts often exhibit reversible formation, leading to equilibria in solution characterized by formation constants (K_eq) that reflect their stability. For example, triarylborane-amine adducts display K_eq values spanning several orders of magnitude, indicating tunable stability based on steric and electronic factors. Bond lengths in coordinate bonds are typically slightly longer than in analogous covalent bonds; the N-B distance in H₃N→BF₃ is about 1.60 Å, compared to 1.47 Å for a standard N-C single bond. Thermodynamically, adduct formation is generally exothermic in the gas phase, but dissociation can be entropy-driven in solution due to increased molecular freedom.

Polarity and Reactivity

The inherent polarity of a coordinate covalent bond arises from the donation of an electron pair from the donor atom to the acceptor atom, resulting in an uneven distribution of electron density. The donor atom, typically more electronegative, retains a partial negative charge (δ⁻), while the acceptor atom develops a partial positive charge (δ⁺), creating a significant dipole moment across the bond. This charge separation is evident in molecules like ammonia-borane (NH₃·BH₃), where the nitrogen donor exhibits higher electron density compared to the boron acceptor. In the ammonium ion (NH₄⁺), the four coordinate bonds from nitrogen to hydrogen further illustrate this, with nitrogen bearing δ⁻ and each hydrogen δ⁺, enhancing the overall molecular dipole. This polarity profoundly influences reactivity by imparting enhanced nucleophilicity to the donor site and electrophilicity to the acceptor site. The partial negative charge on the donor atom increases its tendency to participate in electrophilic additions or further bonding, while the positive charge on the acceptor makes it prone to nucleophilic attacks. In coordination complexes, such as those of transition metals, the electrophilic metal center (acceptor) readily undergoes ligand exchange, where an incoming nucleophilic ligand displaces the existing one through associative or dissociative mechanisms. This reactivity is driven by the bond's polar nature, which lowers the activation barrier for electron transfer processes compared to symmetric bonds. Spectroscopic techniques provide direct evidence of this polarity through shifts in vibrational modes. Infrared (IR) spectroscopy reveals changes in bond stretching frequencies upon coordinate bond formation; for example, the N-H stretching vibrations in NH₄⁺ occur at lower frequencies (around 3000–3200 cm⁻¹) and broader bands compared to free NH₃ (3330–3450 cm⁻¹), due to the increased s-character and electrostatic effects from the polar coordinate bonds. These shifts confirm the altered electron density and bond strength induced by the dative interaction. The polar character also affects solvation, promoting stronger interactions with polar solvents that stabilize the charge-separated structure through dipole-dipole or ion-dipole forces. Computational and experimental studies demonstrate that dative bond stability increases markedly in more polar solvents, as the solvent dielectric screens the partial charges and reduces destabilizing interactions. In contrast to non-polar covalent bonds, which cleave homolytically to form radicals, coordinate bonds favor heterolytic cleavage, where the electron pair remains with the donor atom, aligning with the bond's polarity direction and facilitating ionic reaction pathways.

Examples

Simple Molecules

One prominent example of a coordinate covalent bond occurs in the ion, [NH₄]⁺, formed when (NH₃) donates its of electrons on the atom to a proton (H⁺). In the , three N–H bonds are conventional covalent bonds, while the fourth is a dative bond represented by an arrow from N to H, though all four bonds become equivalent in the tetrahedral geometry of the ion. Similarly, the oxonium ion, [H₃O]⁺, arises from the donation of a from (H₂O) to H⁺. The Lewis structure shows two O–H covalent bonds and one dative O→H bond, resulting in a trigonal pyramidal where all three bonds are indistinguishable. In the boron trifluoride-ammonia adduct, H₃N→BF₃, ammonia acts as the donor, providing its nitrogen lone pair to the electron-deficient boron in BF₃. The Lewis structure depicts the dative bond as an arrow from N to B, with boron changing from trigonal planar geometry in BF₃ (due to sp² hybridization) to tetrahedral in the adduct (sp³ hybridization), altering the B–F bond angles and lengths.

Coordination Complexes

In coordination chemistry, coordinate covalent bonds, also known as dative bonds, form the foundation of metal-ligand interactions where ligands donate electron pairs to the central metal ion. Neutral ligands such as ammonia (NH₃) provide lone pairs from their nitrogen atoms, while anionic ligands like chloride (Cl⁻) donate from their filled orbitals, resulting in bonds denoted as M←L, with the arrow indicating the direction of electron pair donation from ligand to metal. These interactions position the metal as a Lewis acid accepting the electron density, enabling the assembly of stable complexes. A classic example is the hexaamminecobalt(III) ion, [Co(NH₃)₆]³⁺, where the Co³⁺ ion is surrounded by six NH₃ ligands, each forming a dative N→Co bond through the nitrogen lone pair, leading to an overall charge of +3 on the complex. In biological systems, similar bonds occur in hemoglobin, where the iron(II) center forms Fe←N dative bonds with nitrogen atoms from the porphyrin ring and a proximal histidine residue, facilitating oxygen transport. The coordination number, defined as the number of such dative bonds to the metal, dictates the geometry; for instance, a coordination number of six often yields octahedral arrangements as in [Co(NH₃)₆]³⁺, while four-coordinate complexes like those of d⁸ metals such as Pt(II) adopt square planar geometries influenced by the directional nature of these bonds. Back-bonding further modulates these interactions in many complexes, where the metal donates electrons from its filled d-orbitals into the empty π* antibonding orbitals of the ligand, enhancing overall bond strength through synergistic σ-donation and π-acceptance. This π-backbonding is particularly prominent in complexes with ligands like CO or phosphines, stabilizing low-oxidation-state metals. In nomenclature, dative bonds are conventionally represented with arrows in structural formulas, such as ligand → metal, to distinguish the electron origin, though IUPAC names typically omit this notation in favor of standard ligand descriptors.

Comparisons

With Covalent Bonds

A coordinate , also known as a dative bond, differs from a standard primarily in the origin of the shared . In a regular , each participating atom contributes one electron to the shared pair, resulting in mutual sharing. In contrast, a coordinate forms when a single atom donates both electrons from a lone pair to an electron-deficient atom, such as a cation or species with an empty orbital. Once established, however, a coordinate covalent bond becomes indistinguishable from a regular covalent bond in terms of strength, stability, and overall properties. The shared electrons behave equivalently in the final molecular structure, with no experimental method able to differentiate the bond type based on physical characteristics like bond length or dissociation energy. This equivalence is illustrated in the ammonium ion, [\ceNH4+][\ce{NH4+}], where three N-H bonds form through standard covalent sharing in ammonia (\ceNH3\ce{NH3}), while the fourth involves nitrogen donating its lone pair to a proton (\ceH+\ce{H+}). Despite the distinct formation pathways—a regular C-H bond in methane (\ceCH4\ce{CH4}) involves equal electron contribution from carbon and hydrogen—the N-H bonds in [\ceNH4+][\ce{NH4+}] are symmetric and identical in nature post-formation. Both bond types rely on orbital overlap for formation, typically involving sigma bonds from the end-on interaction of atomic orbitals, such as s or p orbitals. In coordinate bonds, the donor atom's filled orbital overlaps with the acceptor atom's empty orbital, similar to regular covalent overlap, but the process often leads to greater polarity due to the initial charge imbalance from the lone pair donation. This can result in partial charges, as seen in adducts like \ceNH3⋅BF3\ce{NH3 \cdot BF3}, where the bond creates formal positive and negative charges on nitrogen and boron, respectively, enhancing the bond's dipole moment compared to many symmetric covalent bonds. The distinction between coordinate and regular covalent bonds is most relevant during the initial formation stage, influencing reaction kinetics in Lewis acid-base interactions, and in representational tools like Lewis structures, where coordinate bonds are denoted by arrows to indicate electron donation direction. For instance, the stepwise reaction \ceNH3+H+−>NH4+\ce{NH3 + H+ -> NH4+} highlights the kinetic role of the lone pair donation, which would not apply to symmetric covalent formations.

With Ionic Bonds

Ionic bonds form through the complete transfer of one or more valence electrons from a metal atom to a non-metal atom, creating positively charged cations and negatively charged anions that are attracted to each other by electrostatic forces. This electron transfer results in discrete ions rather than shared electron pairs, distinguishing ionic bonding from all covalent variants. Coordinate covalent bonds, as a subtype of covalent bonding, involve the sharing of an electron pair where both electrons originate from one atom (the donor), but the pair is shared between both atoms without full transfer. Although these bonds can exhibit significant polarity due to the unequal contribution and electronegativity differences, the electrons remain shared, preventing the formation of separate ions as in ionic bonds. Coordinate bonds occupy a position in the spectrum of chemical bonding, bridging purely ionic and purely covalent types, particularly through mechanisms described by Fajans' rules. These rules state that a bond's character shifts toward covalent when a small, highly charged cation polarizes a large, polarizable anion, distorting the electron cloud and inducing partial sharing. In such cases, what begins as an ionic interaction can develop covalent features, including coordinate sharing in adducts or complexes. A representative example of an ionic bond is sodium chloride (NaCl), where sodium transfers an electron to chlorine, forming Na⁺ and Cl⁻ ions held by electrostatic attraction. In comparison, the tetrachloroaluminate anion ([AlCl₄]⁻) illustrates a coordinate covalent bond, formed when aluminum chloride (AlCl₃) accepts a lone pair from a chloride ion (Cl⁻), resulting in all four Al-Cl bonds being equivalent with shared electrons. Regarding physical properties, ionic compounds like NaCl typically exhibit high solubility in polar solvents such as water and conduct electricity in molten or aqueous states due to mobile ions. Coordination compounds containing coordinate covalent bonds vary in behavior: neutral molecular complexes often act as discrete entities with lower solubility in water and minimal conductivity, as they do not dissociate into ions. However, many coordination compounds exist as ionic salts, such as [Co(NH₃)₆]Cl₃, which dissociate into complex ions and counterions in solution, exhibiting high solubility and electrical conductivity similar to ionic compounds.

References

  1. https://goldbook.iupac.org/terms/view/C01329
  2. https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_%28Physical_and_Theoretical_Chemistry%29/Chemical_Bonding/Fundamentals_of_Chemical_Bonding/Coordinate_%28Dative_Covalent%29_Bonding
  3. https://www.geo.utexas.edu/courses/376m/coord_chem.htm
  4. https://bouman.chem.georgetown.edu/S02/lect31/lect31.htm
  5. https://goldbook.iupac.org/terms/view/D01523
  6. https://www2.chemistry.msu.edu/courses/cem151/chap24lect_2024.pdf
  7. https://chem.libretexts.org/Courses/University_of_Kentucky/UK%253A_General_Chemistry/15%253A_Equilibria_of_Other_Reaction_Classes/15.2%253A_Lewis_Acids_and_Bases
  8. https://openstax.org/books/chemistry-2e/pages/15-2-lewis-acids-and-bases
  9. https://courses.lumenlearning.com/suny-chem-atoms-first/chapter/lewis-acids-and-bases-2/
  10. https://chem-textbook.ucalgary.ca/version2/chapter-15-main/lewis-acids-and-bases/
  11. https://chem.libretexts.org/Courses/Smith_College/Organic_Chemistry_%28LibreTexts%29/02%253A_Polar_Covalent_Bonds_Acids_and_Bases/2.11%253A_Acids_and_Bases_-_The_Lewis_Definition
  12. https://byjus.com/chemistry/lewis-acids-and-bases/
  13. https://pubs.acs.org/doi/pdf/10.1021/cr60074a005
  14. https://www.nobelprize.org/prizes/chemistry/1913/werner/facts/
  15. https://www.sciencehistory.org/education/scientific-biographies/gilbert-newton-lewis/
  16. https://onlinelibrary.wiley.com/doi/full/10.1002/jcc.20517
  17. https://www.britannica.com/biography/Nevil-Vincent-Sidgwick
  18. https://www.britannica.com/biography/Gilbert-N-Lewis/Chemical-bonding-theory
  19. https://chem.libretexts.org/Bookshelves/Inorganic_Chemistry/Introduction_to_Inorganic_Chemistry_(Wikibook)/03:_Acid-Base_Chemistry/3.02:_Brnsted_and_Lewis_Acids_and_Bases
  20. https://open.maricopa.edu/chemistryfundamentals/chapter/lewis-acids-and-bases-2/
  21. https://users.highland.edu/~jsullivan/principles-of-general-chemistry-v1.0/s12-07-lewis-acids-and-bases.html
  22. https://www.utdallas.edu/~son051000/chem1312/Chapter16a.pdf
  23. https://www.mdpi.com/1422-0067/5/8/239
  24. https://pubs.rsc.org/en/content/articlelanding/2010/cc/c0cc00408a
  25. https://pubs.acs.org/doi/10.1021/ja01186a085
  26. https://pubs.acs.org/doi/10.1021/acs.jpca.1c05766
  27. https://pubs.rsc.org/en/content/articlepdf/1994/c3/c39940001489
  28. https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Chemical_Bonding/Fundamentals_of_Chemical_Bonding/Coordinate_(Dative_Covalent)_Bonding
  29. https://chem.libretexts.org/Bookshelves/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Coordination_Chemistry/Structure_and_Nomenclature_of_Coordination_Compounds/Ligand_Field_Theory/Ligand_Substitution_Reactions
  30. https://pubs.acs.org/doi/10.1021/j100159a020
  31. https://www.nature.com/articles/s41467-022-29806-3
  32. https://www.pharmacy180.com/article/heterolytic-bond-cleavages-1462/
  33. https://www.chem.purdue.edu/gchelp/cchem/
  34. https://chemed.chem.purdue.edu/genchem/topicreview/bp/ch12/complex.php
  35. https://earth.callutheran.edu/Academic_Programs/Departments/BioDev/omm/jsmol/hemo/hemoglobina.html
  36. https://nwclug.harpercollege.edu/libweb/DBLQNn/272013/CoordinationChemistry.pdf
  37. https://alpha.chem.umb.edu/chemistry/ch611/documents/Lec5-pibonding.pdf
  38. https://archive.cbts.edu/Textbook/16Mso3/418168/Coordinate-Bond-And-Covalent-Bond.pdf
  39. https://www2.chem.wisc.edu/deptfiles/genchem/sstutorial/Text7/Tx71/tx71.html
  40. https://chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_Chemistry_-_The_Central_Science_(Brown_et_al.)/08%3A_Basic_Concepts_of_Chemical_Bonding/8.02%3A_Ionic_Bonds
  41. https://chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_Chemistry_-_The_Central_Science_(Brown_et_al.)/09%3A_Covalent_Bonding/9.01%3A_Bonding_and_Molecules
  42. https://chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_Chemistry_-_The_Central_Science_(Brown_et_al.)/08%3A_Basic_Concepts_of_Chemical_Bonding/8.04%3A_Polarity_and_the_Electronegativity_Scale
  43. https://acshist.scs.illinois.edu/awards/OPA%20Papers/1990-Holmen1.pdf
  44. https://chem.libretexts.org/Bookshelves/Inorganic_Chemistry/Supplemental_Modules_%28Inorganic_Chemistry%29/Coordination_Chemistry/Structure_and_Nomenclature_of_Coordination_Compounds/Electron_Counting
  45. https://chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_Chemistry_-_The_Central_Science_(Brown_et_al.)/02%3A_Atoms_Molecules_and_Ions/2.07%3A_Ions_and_Ionic_Compounds
  46. https://chem.libretexts.org/Bookshelves/Inorganic_Chemistry/Supplemental_Modules_and_Websites_%28Inorganic_Chemistry%29/Coordination_Chemistry/Structure_and_Nomenclature_of_Coordination_Compounds/Introduction_to_Coordination_Chemistry
  47. https://chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_Principles_of_Modern_Chemistry_(Oxtoby_et_al.)/Unit_2%3A_Atoms_and_Molecules/08%3A_Covalent_Bonding_and_Lewis_Structures/8.05%3A_Resonance_and_Formal_Charge
Add your contribution
Related Hubs
User Avatar
No comments yet.