Silicon tetrachloride
Silicon tetrachloride is an inorganic compound with the chemical formula SiCl₄, consisting of a silicon atom bonded to four chlorine atoms in a tetrahedral geometry.[1] It appears as a colorless, fuming liquid at room temperature with a pungent odor, boiling at 57.6 °C and decomposing violently with water to form silicic acid and hydrogen chloride gas, releasing significant heat.[1] Industrially produced by the chlorination of silicon or silica in the presence of carbon, it serves primarily as a key precursor for high-purity polycrystalline silicon via the Siemens process, essential for manufacturing semiconductor wafers and photovoltaic cells.[2][3] Due to its reactivity, silicon tetrachloride is highly corrosive, causing severe burns to skin and eyes upon contact, and requires stringent handling precautions including protective equipment and avoidance of moisture.[4] Its applications extend beyond electronics to optical fiber production and as a source of silicon in chemical vapor deposition processes, underscoring its role in advanced materials synthesis.[5]History
Discovery and early production
Silicon tetrachloride (SiCl₄) was first prepared in 1823 by the Swedish chemist Jöns Jacob Berzelius, who synthesized it by heating freshly isolated amorphous silicon in a stream of chlorine gas at elevated temperatures.[6][7] Berzelius had obtained the silicon precursor earlier that year through the reduction of potassium hexafluorosilicate (K₂SiF₆) with metallic potassium, yielding an impure but elemental form of silicon that reacted vigorously with Cl₂ to form the volatile, colorless liquid SiCl₄.[8] This direct combination, Si + 2 Cl₂ → SiCl₄, represented the inaugural empirical demonstration of silicon's tetravalent halide chemistry, highlighting its parallels to carbon tetrachloride in forming stable, tetrahedral structures.[9] Early production remained confined to small-scale laboratory efforts, as pure silicon sources were scarce and the reaction required careful control to avoid side products from impurities in the silicon. Berzelius's method underscored silicon's high reactivity with halogens, predating broader recognition of the element's properties and enabling subsequent reductions of SiCl₄ with metals like sodium or magnesium to yield purer silicon samples in the mid-19th century.[10] These initial syntheses advanced causal understanding of silicon's bonding behavior, establishing SiCl₄ as a key intermediate in early silicon chemistry without reliance on modern purification techniques.[11]Physical and chemical properties
Physical characteristics
Silicon tetrachloride is a colorless, fuming liquid with a pungent odor.[1] Its melting point is −68.9 °C, and its boiling point is 57.6 °C at standard pressure.[12][13] The density is 1.48 g/cm³ at 20 °C.[12][14] It exhibits high vapor pressure, approximately 26 kPa at 20 °C, which causes it to fume in contact with moist air owing to rapid hydrolysis.[14] Silicon tetrachloride reacts exothermically with water rather than dissolving, producing hydrochloric acid and silicic acid; it is miscible with organic solvents including benzene, toluene, chloroform, and ether.[1][14]| Property | Value | Conditions |
|---|---|---|
| Appearance | Colorless liquid | Room temperature |
| Odor | Pungent | - |
| Molar mass | 169.90 g/mol | - |
| Vapor pressure | 26 kPa | 20 °C |
Reactivity and bonding
Silicon tetrachloride exhibits tetrahedral coordination at the central silicon atom, achieved through sp³ hybridization of its valence orbitals, which form four equivalent sigma bonds with chlorine atoms. This hybridization arises from the mixing of silicon's 3s and three 3p orbitals, accommodating the four bonding pairs while adhering to the octet rule in the ground state molecule.[15][16] The silicon center's reactivity stems from its inherent Lewis acidity, driven by the electronegativity difference between silicon (1.90) and chlorine (3.16), resulting in polar Si–Cl bonds with partial positive charge on silicon. This electron deficiency enables SiCl₄ to accept electron pairs from Lewis bases, forming adducts such as SiCl₄·NH₃, and promotes nucleophilic cleavage of Si–Cl bonds. Availability of empty 3d orbitals on silicon further facilitates hypervalency in transition states or intermediates, allowing coordination numbers beyond four during interactions.[17] In protic environments, SiCl₄ displays thermodynamic instability due to the greater bond strength of silicon-oxygen interactions compared to silicon-chlorine bonds, rendering hydrolysis highly exergonic and exothermic. This drives spontaneous reaction with water, evolving substantial heat and underscoring the compound's sensitivity to nucleophilic protic species.[18][19]Synthesis
Industrial methods
Silicon tetrachloride is primarily produced industrially through the carbochlorination of silica sand (SiO₂) with carbon and chlorine gas, following the reaction SiO₂ + 2C + 2Cl₂ → SiCl₄ + 2CO.[20] This process occurs in a high-temperature reactor, typically at 800–1200°C, where silica and carbon are mixed and exposed to a flow of chlorine gas, facilitating the reduction and chlorination to yield SiCl₄ vapor alongside carbon monoxide byproduct.[21] [22] The reaction efficiency depends on factors such as carbon-to-silica ratio, gas flow rates, and residence time, with yields optimized around 70–80% silicon volatilization under controlled conditions.[20] A significant portion of industrial silicon tetrachloride arises as a byproduct during polysilicon manufacturing via the Siemens process, where trichlorosilane (HSiCl₃) undergoes chemical vapor deposition on heated silicon rods, producing silicon deposits and SiCl₄ as a waste stream.[23] [24] In this method, for every mole of silicon deposited, 3–4 moles of SiCl₄ are generated, necessitating recycling or separate purification to mitigate environmental and economic losses.[25] Crude silicon tetrachloride from these processes is purified primarily through fractional distillation to remove volatile impurities like unreacted chlorine, hydrogen-containing silanes, and metal chlorides, achieving semiconductor-grade purity levels exceeding 99.999%.[26] [27] Additional steps, such as adsorption or reactive treatments, may target specific contaminants like hydrogen compounds or organics for electronic applications.[28]Laboratory-scale preparation
Silicon tetrachloride can be prepared in the laboratory by direct chlorination of elemental silicon with chlorine gas. The reaction, Si + 2 Cl₂ → SiCl₄, is conducted by heating silicon powder or granules to 500–700 °C in a quartz or glass tube furnace while passing a stream of dry chlorine gas, allowing the volatile product to be collected by condensation in a cooled trap.[29] This method prioritizes purity and control, yielding SiCl₄ suitable for spectroscopic or synthetic applications, though yields depend on silicon particle size and gas flow rates to ensure complete reaction without side products like dichlorodisilane.[30] Ferrosilicon, an alloy containing 75–90% silicon, serves as an accessible alternative starting material for small-scale preparations, reacting similarly under chlorine flow at comparable temperatures to produce SiCl₄ alongside iron chlorides, which can be separated by fractional distillation.[29] The process requires inert atmosphere handling to prevent moisture-induced hydrolysis prior to use. Historically, Jöns Jakob Berzelius first isolated SiCl₄ in 1823–1824 by reacting newly prepared amorphous silicon with chlorine gas, marking the compound's initial synthesis amid early efforts to characterize silicon halides.[9] This direct approach remains a benchmark for laboratory verification of the reaction's kinetics and thermodynamics, often studied under controlled conditions to model vapor-phase chlorination.[31]Molecular structure
Geometric and electronic features
Silicon tetrachloride exhibits a tetrahedral geometry with Td point group symmetry, characterized by four equivalent Si-Cl bonds and Cl-Si-Cl bond angles of 109.471°. The experimental Si-Cl bond length is 2.019 Å, determined via electron diffraction studies. This configuration arises from the sp³ hybridization of the central silicon atom, which accommodates four sigma bonds to chlorine ligands, minimizing electron repulsion according to valence shell electron pair repulsion theory while aligning with quantum mechanical predictions of minimal energy.[32] The electronic structure features sigma bonding molecular orbitals primarily formed by the overlap of silicon 3s and 3p orbitals with chlorine 3p orbitals, yielding a closed-shell singlet ground state. The lowest unoccupied molecular orbital (LUMO), of antibonding σ* character centered on silicon, resides at relatively low energy, enhancing the molecule's electrophilicity and capacity for Lewis acid behavior. This is evidenced by the ionization energy of 11.79 eV, reflecting the stability of the occupied orbitals.[33] Vibrational spectroscopy confirms the Td symmetry through distinct infrared (IR) and Raman-active modes. The fundamental vibrations include the Raman-active symmetric stretch ν₁ (A₁) at 424 cm⁻¹, the Raman-active doubly degenerate deformation ν₂ (E) at 150 cm⁻¹, the IR-active triply degenerate asymmetric stretch ν₃ (T₂) at 621 cm⁻¹, and the triply degenerate deformation ν₄ (T₂) at 221 cm⁻¹, which is active in both IR and Raman spectra. These assignments, derived from gas-phase measurements, underscore the bond strengths and symmetry-forbidden/inactive modes expected for a tetrahedral molecule, such as the IR-inactive ν₁.Comparisons with analogous compounds
Silicon tetrachloride (SiCl₄) shares a tetrahedral molecular geometry with its carbon analog, carbon tetrachloride (CCl₄), but exhibits markedly different reactivity, particularly in hydrolysis. While CCl₄ remains stable and inert toward water even under prolonged exposure, SiCl₄ reacts vigorously to form silicic acid and hydrogen chloride, driven by silicon's larger atomic radius (110 pm vs. 70 pm for carbon), which permits easier access for nucleophilic water molecules despite the steric bulk of the chloride ligands, and the presence of vacant 3d orbitals that allow temporary expansion of the coordination number to accommodate the attacking oxygen.[34][35] Carbon lacks such d-orbitals and has stronger pπ-pπ overlap in potential intermediate states, rendering C–O bond formation thermodynamically unfavorable relative to C–Cl retention. This disparity persists despite the Si–Cl bond dissociation energy being stronger (~381 kJ/mol) than the average C–Cl bond (~327–341 kJ/mol in CCl₄), as the reaction's driving force stems from the high stability of Si–O bonds (452 kJ/mol) over Si–Cl, enabling exothermic hydrolysis overall.[36][37] Comparisons with heavier group 14 homologs reveal trends in thermal stability and physical properties. The tetrachlorides' stability decreases down the group (CCl₄ > SiCl₄ > GeCl₄ > SnCl₄), attributable to increasing metallic character and weaker M–Cl bonding due to poorer orbital overlap with diffuse valence orbitals; for instance, SnCl₄ decomposes to SnCl₂ and Cl₂ upon heating above 150 °C, whereas SiCl₄ withstands temperatures up to ~500 °C without decomposition.[38][39] SiCl₄'s volatility (boiling point 57.6 °C) exceeds that of GeCl₄ (83.1 °C) and the liquid SnCl₄ (114 °C at standard pressure), arising from its lower molecular weight and minimal intermolecular forces in the tetrahedral structure, which facilitates its use in vapor-phase processes like epitaxial deposition, unlike the less volatile heavier analogs.[40]| Compound | Boiling Point (°C) | Thermal Stability |
|---|---|---|
| CCl₄ | 76.7 | High (inert to heat) |
| SiCl₄ | 57.6 | Moderate (stable to ~500 °C) |
| GeCl₄ | 83.1 | Lower than SiCl₄ |
| SnCl₄ | 114 | Low (decomposes >150 °C) |