Group 12 element
Group 12 elements, also known as the zinc group, comprise zinc (Zn, atomic number 30), cadmium (Cd, atomic number 48), mercury (Hg, atomic number 80), and the synthetic superheavy element copernicium (Cn, atomic number 112).[1][2] These d-block elements feature a filled (n-1)d¹⁰ subshell and ns² valence electrons in their neutral state, resulting in a predominant +2 oxidation state and properties that distinguish them from typical transition metals, as they lack partially filled d orbitals in their common ions.[1][3] The lighter members—zinc, cadmium, and mercury—are soft, silvery-white metals with relatively low melting and boiling points compared to other d-block elements, reflecting weak metallic bonding due to the tightly held ns electrons.[1] Zinc has a melting point of 419.53°C and density of 7.14 g/cm³, cadmium melts at 321.07°C with a density of 8.65 g/cm³, and mercury is unique as the only metal liquid at room temperature, with a melting point of -38.83°C and density of 13.534 g/cm³.[1] These elements occur naturally, primarily in sulfide ores such as zinc blende (ZnS) for zinc and cinnabar (HgS) for mercury, with cadmium often obtained as a byproduct of zinc processing; their crustal abundances decrease down the group, from 75 ppm for zinc to 0.05 ppm for mercury.[1][3] Chemically, Group 12 elements exhibit similarities to alkaline earth metals in their +2 compounds, forming ionic halides that become increasingly covalent down the group, and showing a strong affinity for soft ligands like sulfur and phosphorus.[3] Zinc and cadmium react readily with acids and oxygen to form amphoteric oxides, while mercury requires higher temperatures for oxidation and forms the distinctive Hg₂²⁺ ion in +1 compounds.[3] Applications include zinc in galvanizing steel and alloys like brass, cadmium in rechargeable batteries (though increasingly phased out due to toxicity), and mercury in historical thermometers and the chlor-alkali process (phased out globally by 2025), despite environmental concerns over its bioaccumulation.[3][4] Copernicium, synthesized in 1996 by a German team via heavy-ion bombardment, is highly radioactive with isotopes like ²⁸⁵Cn having half-lives of seconds, limiting experimental study to predicted properties based on relativistic quantum calculations.[2] Theoretical predictions suggest copernicium would be a volatile liquid at room temperature with a +2 oxidation state, exhibiting noble gas-like properties due to relativistic effects stabilizing its 7s electrons, deviating from typical group 12 behavior.[5] Overall, the group highlights trends in metallic character and reactivity influenced by increasing atomic size and lanthanide contraction, with practical and toxicological significance in materials science and environmental chemistry.[1]Introduction
Characteristics
The group 12 elements of the periodic table consist of zinc (Zn, atomic number 30), cadmium (Cd, 48), mercury (Hg, 80), and copernicium (Cn, 112).[2] These are d-block metals distinguished by their completely filled d subshell in the ground state, with an electron configuration of noble gasd^{10} ns^2, which imparts properties more characteristic of main-group elements than those of typical transition metals with incomplete d orbitals.[6] The dominant oxidation state across the group is +2, resulting from the facile loss of the ns^2 valence electrons, while access to d electrons for higher oxidation states is limited due to the stable closed-shell configuration.[6] Zinc and cadmium appear as silvery-white metals with a bluish tinge, solid and malleable at room temperature, reflecting their metallic bonding despite the full d subshell weakening interatomic interactions compared to earlier d-block groups.[7][8] Mercury, however, is the only metal that remains liquid at standard conditions, its silvery appearance and fluidity stemming from particularly weak metallic bonding influenced by relativistic stabilization of the 6s electrons.[9] Copernicium, entirely synthetic and produced in trace amounts via nuclear reactions, lacks observed bulk properties, but computational models predict it as a volatile liquid under ambient conditions, with a melting point of approximately 283 K and boiling point of 340 K, arising from strong relativistic effects that contract the 7s orbital and expand the 6d, yielding a density akin to mercury's alongside noble-gas-like insulating behavior marked by a 6.4 eV band gap and dispersion-dominated cohesion.[10] Owing to their d^{10} configuration, group 12 elements exhibit post-transition metal traits, including behaviors reminiscent of p-block elements such as restricted oxidation state variability and a propensity for covalent or amphoteric compound formation rather than the diverse coordination chemistry of true transition metals.[11] They also show a diagonal relationship with group 2 elements, notably zinc and magnesium, where comparable charge-to-radius ratios foster similar reactivities, as seen in the influence of zinc(I) complexes on the development of stable magnesium(I) dimers.[12] This positions group 12 as a bridge between transition and main-group chemistries, with copernicium's relativistic deviations further emphasizing the group's evolving properties down the period.[10]Position in the periodic table
Group 12 elements occupy the 12th column of the periodic table within the d-block, situated between group 11—comprising the coinage metals copper, silver, and gold—and group 13, which initiates the post-transition metals. This placement aligns them with other d-block elements in groups 3 through 12, where the (n-1)d subshell is progressively filled across the periods.[13][14] The general electronic configuration of these elements is ns^2 (n-1)d^{10}, characterized by a completely filled d subshell and two valence s electrons, which distinguishes their electronic structure from earlier d-block groups that exhibit partially filled d orbitals. This configuration contributes to their reactivity patterns, primarily involving the loss of the ns electrons to form +2 oxidation states, while the d electrons remain largely uninvolved in bonding. For the heavier members, particularly mercury, the inert pair effect begins to manifest, stabilizing the ns² electrons and reducing their participation in chemical bonds compared to lighter analogs like zinc.[13][14] Due to the full d¹⁰ subshell in both neutral atoms and common ions (e.g., Zn²⁺ with [Ar] 3d¹⁰), group 12 elements lack the incomplete d subshell required by the IUPAC definition of transition metals, leading to their frequent classification alongside main-group metals despite their d-block position. This absence of d-orbital involvement in bonding results in properties more akin to post-transition metals, such as lower melting points and limited variable oxidation states, setting them apart from groups 3–11.[14]/Descriptive_Chemistry/Elements_Organized_by_Block/3_d-Block_Elements/1b_Properties_of_Transition_Metals/General_Trends_among_the_Transition_Metals) In mercury and the superheavy copernicium, relativistic effects significantly alter these trends through s-orbital contraction, where high nuclear charge accelerates inner electrons, increasing their effective mass and shrinking the 6s (Hg) or 7s (Cn) orbitals, thereby elevating ionization energies (e.g., Hg 6s binding energy of 8.92 eV relativistically vs. 7.10 eV non-relativistically). This contraction stabilizes the valence s electrons, enhancing the inert pair effect and reducing metallic bonding strength in mercury, contributing to its liquid state at room temperature. For copernicium, these effects are amplified by its superheavy nature, predicting a highly volatile noble-liquid behavior with a melting point around 10°C and boiling point near 67°C, alongside potential d-character in bonding that deviates from typical group 12 trends due to nuclear instability and short isotope half-lives (up to 29 seconds).[15][16][5]Physical properties
Atomic properties
The group 12 elements—zinc (Zn), cadmium (Cd), mercury (Hg), and the synthetic copernicium (Cn)—exhibit a common valence electron configuration of ns^2 (n-1)d^{10}, where the filled d subshell contributes to their relatively stable +2 oxidation state, though relativistic effects become prominent in the heavier members.[17] This configuration arises from the addition of two s electrons outside a closed d shell, distinguishing them from typical transition metals with partially filled d orbitals. The atomic masses and full electron configurations are summarized below, with values for Cn being theoretical predictions based on relativistic quantum calculations.| Element | Atomic Number | Atomic Mass (u) | Electron Configuration |
|---|---|---|---|
| Zinc (Zn) | 30 | 65.38 | [ \ce{Ar} ] 3d^{10} 4s^2 |
| Cadmium (Cd) | 48 | 112.41 | [ \ce{Kr} ] 4d^{10} 5s^2 |
| Mercury (Hg) | 80 | 200.59 | [ \ce{Xe} ] 4f^{14} 5d^{10} 6s^2 |
| Copernicium (Cn) | 112 | 285 (predicted) | [ \ce{Rn} ] 5f^{14} 6d^{10} 7s^2 (predicted) |
Bulk properties
Group 12 elements exhibit a range of bulk physical properties that reflect their position in the periodic table, with zinc and cadmium behaving as typical metals and mercury as the only metallic element that is liquid at standard temperature and pressure. These properties include phase transition temperatures, densities, conductivities, and crystal structures, which vary significantly down the group due to increasing atomic size and relativistic effects in heavier members. Copernicium, the synthetic superheavy element, has properties predicted through computational methods, indicating it may differ markedly from its homologues. Predictions for copernicium are based on relativistic density functional theory (DFT) calculations due to its short half-life preventing direct measurement. The melting and boiling points decrease down the group, highlighting the trend toward weaker metallic bonding in heavier elements. Zinc has a melting point of 419.5 °C and a boiling point of 907 °C, while cadmium melts at 321.1 °C and boils at 767 °C; mercury, in contrast, melts at -38.8 °C and boils at 356.7 °C.[24][25] Densities increase with atomic mass, from 7.14 g/cm³ for zinc to 8.65 g/cm³ for cadmium and 13.53 g/cm³ for mercury.[26] For copernicium, density functional theory (DFT) calculations predict a density of approximately 13.7 g/cm³, similar to mercury.[27]| Element | Melting Point (°C) | Boiling Point (°C) | Density (g/cm³) |
|---|---|---|---|
| Zinc | 419.5 | 907 | 7.14 |
| Cadmium | 321.1 | 767 | 8.65 |
| Mercury | -38.8 | 356.7 | 13.53 |
| Copernicium (predicted) | ~10 | ~67 | ~13.7 |
Chemical properties
Periodic trends
The elements of group 12—zinc (Zn), cadmium (Cd), mercury (Hg), and copernicium (Cn)—exhibit periodic trends in their physical and chemical properties that deviate from typical d-block behavior due to their filled d¹⁰ electronic configurations and increasing relativistic effects down the group. Reactivity decreases from Zn to Hg, as evidenced by the standard reduction potentials for the M²⁺/M couples, which become less negative: Zn²⁺ + 2e⁻ → Zn (E° = -0.76 V), Cd²⁺ + 2e⁻ → Cd (E° = -0.40 V), and Hg²⁺ + 2e⁻ → Hg (E° = +0.85 V).[33] This trend reflects the increasing nobility of the metals, with Zn being the most reactive and readily oxidized, while Hg is resistant to oxidation under standard conditions. For Cn, theoretical predictions suggest even lower reactivity, potentially resembling a noble metal due to enhanced relativistic stabilization of the 7s orbitals. Experimental gas-phase studies, such as adsorption on gold surfaces in 2016, indicate Cn is more noble than Hg, aligning with predictions (as of 2024).[34] A key anomaly in atomic properties arises in the first ionization energies (IE₁), which do not monotonically decrease down the group as expected from increasing atomic size. Zn has IE₁ = 9.39 eV, Cd = 8.99 eV, but Hg = 10.44 eV, higher than Zn's due to relativistic effects that contract the 6s orbital and increase effective nuclear charge on the valence electrons.[35] These effects, including spin-orbit coupling and the Darwin term, become pronounced for Hg, stabilizing the closed-shell Hg²⁺ ion and contributing to its low reactivity.[15] For Cn, relativistic influences are expected to be even stronger, yielding an IE₁ of approximately 12 eV, higher than Hg's and further reducing reactivity. Oxidation states in group 12 are predominantly +2, consistent with the ns² configuration, though +1 is rare and observed mainly in Hg as the Hg₂²⁺ dimer. No +3 state is common across the group, as it would require promotion from the stable d¹⁰ subshell. For Cn, calculations predict +2 as the most stable, with +1 possible but unlikely, and +4 potentially accessible under specific conditions due to 7p involvement, though not observed experimentally.[34] Electronegativity on the Pauling scale increases slightly from Zn (1.65) to Cd (1.69) and more notably to Hg (2.00), reflecting the relativistic contraction that enhances Hg's attraction for electrons.[36] Cn's electronegativity is predicted to be in the range of 1.5–2.0, likely closer to Zn and Cd due to balancing relativistic and lanthanide contraction effects in the 7th period.[37] The stability of group 12 hydrides decreases dramatically down the group, highlighting trends in M–H bonding strength. ZnH₂ decomposes slowly at room temperature and rapidly above 90 °C, forming polymeric structures; CdH₂ decomposes above -20 °C; and HgH₂ is highly explosive, detonating upon slight warming or shock. This trend correlates with weakening M–H bonds from relativistic stabilization of the metal's s electrons, reducing their availability for hydride formation. For CnH₂, theoretical models suggest even lower stability, potentially existing only in matrix isolation.[38]| Property | Zn | Cd | Hg | Cn (predicted) |
|---|---|---|---|---|
| First Ionization Energy (eV) | 9.39 | 8.99 | 10.44 | ~12 |
| Electronegativity (Pauling) | 1.65 | 1.69 | 2.00 | ~1.5–2.0 |
| E° (M²⁺/M) (V) | -0.76 | -0.40 | +0.85 | <<0 |
Classification and bonding
Group 12 elements—zinc (Zn), cadmium (Cd), mercury (Hg), and copernicium (Cn)—are classified as post-transition metals, despite their position in the d-block of the periodic table. This designation arises from their filled d¹⁰ electronic configuration in the +2 oxidation state, which renders the d electrons inert and unavailable for bonding, leading to chemical behavior more akin to p-block main-group metals than typical transition metals with variable oxidation states and d-orbital involvement. Lighter members like Zn and Cd exhibit predominantly metallic character, while heavier ones, particularly Hg, display increasing covalent tendencies due to relativistic effects stabilizing the ns orbitals and enhancing s-electron participation in bonds.[39] The bonding in group 12 elements varies with the nature of the ligands and the element involved. In lighter members, compounds such as zinc oxide (ZnO) and zinc hydroxide (Zn(OH)₂) feature predominantly ionic bonding, characterized by high lattice energies and the transfer of electrons from the metal to oxygen, though with some covalent character due to the moderate electronegativity difference. In contrast, organometallic compounds across the group, such as dimethylzinc ((CH₃)₂Zn) and dimethylmercury ((CH₃)₂Hg), exhibit covalent M–C bonds, where electron sharing dominates owing to the low polarity and sp³-hybridized carbon atoms. Mercury compounds, including HgCl₂ and HgS, further emphasize covalent bonding, influenced by the inert-pair effect and relativistic contraction of the 6s orbital, which promotes weaker, more directional bonds over ionic lattices.[40] In coordination chemistry, group 12 elements in the +2 oxidation state prefer tetrahedral geometries for four-coordinate complexes, a consequence of their d¹⁰ configuration, which lacks crystal field stabilization energy preferences that favor octahedral or square planar arrangements in other transition metals. For instance, the tetrachlorozincate ion, [ZnCl₄]²⁻, adopts a tetrahedral structure to minimize steric repulsion among ligands, as the closed-shell d orbitals impose no directional bonding constraints. This geometric preference extends to Cd and Hg analogs, such as [CdCl₄]²⁻ and [HgI₄]²⁻, reinforcing the post-transition metal character. Unlike typical transition metals, group 12 elements lack variable oxidation states beyond +2 (and rarely +1 for Hg), as the stable d¹⁰ ns² ground state and high ionization energies prevent access to higher or d-involving states.[41][42] Theoretical studies on copernicium highlight its divergence from lighter homologs, predicting enhanced volatility and noble gas-like inertness due to extreme relativistic effects stabilizing the 7s² valence shell and weakening metallic bonding. First-principles calculations indicate Cn exists as a volatile liquid near room temperature, with a predicted melting point of approximately 283 K and boiling point of 340 K, and minimal reactivity toward halogens or chalcogens, resembling radon more than mercury in its low cohesion and diatomic tendencies in the gas phase. These predictions stem from free-energy simulations accounting for spin-orbit coupling, underscoring Cn's potential as a relativistic noble liquid rather than a conventional post-transition metal. Experimental gas-phase studies, such as adsorption on gold surfaces in 2016, indicate Cn is more noble than Hg, aligning with predictions (as of 2024).[27]Comparison with group 2 elements
Group 12 elements, zinc (Zn), cadmium (Cd), and mercury (Hg), exhibit notable similarities to the group 2 alkaline earth metals (Be, Mg, Ca, Sr, Ba) primarily due to their shared +2 oxidation state and tendency to form divalent compounds. Both groups predominantly achieve the +2 oxidation state by losing their ns² valence electrons, leading to analogous ionic species like Zn²⁺ and Mg²⁺ or Cd²⁺ and Ca²⁺. For instance, oxides such as ZnO and MgO are both refractory solids used in similar applications, reflecting comparable ionic bonding in these compounds.[13][43] A key similarity arises from a diagonal relationship between group 12 and group 2 elements, driven by comparable ionic radii and charge densities for their +2 cations. The ionic radius of Zn²⁺ (74 pm) is close to that of Mg²⁺ (72 pm), while Cd²⁺ (95 pm) resembles Ca²⁺ (100 pm), and Hg²⁺ (102 pm) aligns with Sr²⁺ (118 pm) or Ba²⁺ (135 pm), resulting in similar solubilities and reactivities for compounds like sulfates and carbonates. This proximity in size leads to parallel trends in lattice energies and precipitation behaviors, such as the decreasing solubility of sulfates down both groups.[13] Despite these parallels, group 12 elements display greater covalent character in their compounds compared to the more ionic nature of group 2. Higher electronegativities in group 12—Zn (1.65), Cd (1.69), Hg (2.00) versus Mg (1.31), Ca (1.00), Sr (0.95)—promote polarization of bonds, enhancing covalency, particularly in Hg compounds. This is evident in the amphoteric behavior of group 12 hydroxides and oxides, unlike the predominantly basic group 2 counterparts; for example, Zn(OH)₂ dissolves in both acids and strong bases to form [Zn(OH)₄]²⁻, while Mg(OH)₂ is insoluble in bases.[13] The d¹⁰ electron configuration in group 12 contributes to this distinction by increasing effective nuclear charge and poor shielding, which heightens polarization despite similar charge densities to group 2 ions. Consequently, group 12 hydrides, such as ZnH₂, are less stable and more covalent than the ionic hydrides of group 2 (e.g., MgH₂), decomposing readily at lower temperatures. Similarly, group 12 oxides like ZnO exhibit amphoterism, reacting with acids and bases, whereas group 2 oxides like MgO remain basic. These differences underscore the transitional position of group 12, blending main-group ionic traits with d-block covalency.[13][43]Characteristic compounds
Group 12 elements form a variety of characteristic compounds that highlight their transition from metallic to more covalent and volatile behavior down the group. The oxides, halides, sulfides, and organometallic derivatives exemplify these trends, with properties influenced by the d¹⁰ electronic configuration and increasing relativistic effects in heavier members. The oxides of zinc, cadmium, and mercury display acid-base properties that shift from amphoteric to basic to unstable. Zinc oxide (ZnO) is amphoteric, dissolving in acids to form salts like zinc chloride and in strong bases to produce zincates, such as Na₂ZnO₂. It forms via the combustion of zinc:$2\mathrm{Zn} + \mathrm{O_2} \rightarrow 2\mathrm{ZnO}
Cadmium oxide (CdO) exhibits basic character, reacting with acids to yield cadmium salts but showing limited solubility in bases. Mercury(II) oxide (HgO), a red or yellow solid, is thermally unstable and decomposes readily upon heating above 500 °C:
$2\mathrm{HgO} \rightarrow 2\mathrm{Hg} + \mathrm{O_2}
This endothermic decomposition was historically significant in the isolation of oxygen gas.[13] Halides of group 12 elements demonstrate increasing covalent character and Lewis acidity variations. Zinc chloride (ZnCl₂), a white hygroscopic solid, acts as a Lewis acid by accepting electron pairs from donors, forming tetrahedral complexes like [ZnCl₄]²⁻; it catalyzes reactions such as the Fischer esterification due to this property. In contrast, mercury(II) chloride (HgCl₂) is a covalent molecular compound with a linear Cl–Hg–Cl geometry, arising from the sp hybridization of the d¹⁰ Hg²⁺ center; it hydrolyzes slowly in water and is noted for its toxicity.[13] Sulfides represent key minerals and industrial materials for lighter group 12 elements. Zinc sulfide (ZnS) occurs in two polymorphs: sphalerite (cubic zinc blende structure) and wurtzite (hexagonal), both featuring tetrahedral Zn–S coordination; it serves as a white pigment in paints and a phosphor in luminescent applications due to its wide band gap. Cadmium sulfide (CdS) forms a hexagonal wurtzite lattice and is prized as a bright yellow pigment (cadmium yellow) in artists' oils and ceramics, though its use has declined due to toxicity concerns.[13] Organometallic compounds of group 12 elements are highly reactive, with volatility and sensitivity increasing down the group. Dialkylzinc reagents, such as dimethylzinc (Zn(CH₃)₂), are pyrophoric and air-sensitive, undergoing rapid exothermic reactions with oxygen or water to form zinc hydroxide and hydrocarbons; they are employed in asymmetric synthesis and polymerizations as nucleophilic alkylating agents. Alkylmercury derivatives, exemplified by dimethylmercury ((CH₃)₂Hg), exhibit strong covalent bonding and extreme toxicity, readily penetrating skin and bioaccumulating to cause severe neurological damage via inhibition of enzymes like methionine synthase.[13] For copernicium (Cn), the heaviest group 12 element, no compounds have been synthesized experimentally, but relativistic quantum chemical calculations predict noble-gas-like behavior due to strong scalar-relativistic stabilization of the 7s² orbitals. Copernicium monoxide (CnO) is expected to be volatile, with properties akin to a weakly bound diatomic species rather than a stable solid oxide. In contrast, copernicium difluoride (CnF₂) is predicted to be thermodynamically stable, with a formation energy of -2.5 eV relative to atomic Cn and F₂, owing to relativistic enhancement of Cn–F bonding; higher fluorides like CnF₄ may also form under extreme conditions. These predictions, from 2019 models incorporating spin-orbit coupling, remain unverified but highlight Cn's deviation from mercury-like chemistry. Experimental gas-phase studies, such as adsorption on gold surfaces in 2016, indicate Cn is more noble than Hg, aligning with predictions (as of 2024).[10]