User:Benjah-bmm27/Si-O

In chemistry, the silicon–oxygen bond (Si\sO bond) is a chemical bond between silicon and oxygen atoms that can be found in many inorganic and organic compounds. In a silicon–oxygen bond, electrons are shared between the two atoms but unequally, with oxygen taking the larger share. Silicon–oxygen bonds are therefore described as covalent and polar, with a partial positive charge on silicon and a partial negative charge on oxygen: Siδ+&mdash;Oδ&minus;. Silicon–oxygen single bonds are longer (1.6 vs 1.4 Å) but stronger (452 vs. about 360 kJ mol−1) than carbon–oxygen single bonds. However, silicon–oxygen double bonds are weaker than carbon–oxygen double bonds (715 vs. 590 kJ mol−1). For these reasons, carbon dioxide is a molecular gas containing two C=O double bonds per carbon atom whereas silicon dioxide is a polymeric solid containing four Si–O single bonds per silicon atom; molecular SiO2 containing two Si=O double bonds would polymerise. Compounds containing silicon–oxygen bonds include materials of major geological and industrial significance such as silica, silicate minerals and silicone polymers like polydimethylsiloxane.

Polarity
In a silicon–oxygen bond, electrons are shared between the two atoms but unequally, with oxygen taking the larger share due to its greater electronegativity. Silicon–oxygen bonds are therefore described as covalent and polar.

Silicon and oxygen atoms share their valence electrons with each other, so silicon–oxygen bonds are described as covalent. However, silicon less electronegative than oxygen so the electron density in the bond is skewed towards oxygen and the bond is polar. The unequal sharing of the electrons in the bond leaves a partial positive charge on silicon and a partial negative charge on oxygen, (Siδ+&mdash;Oδ&minus;).

On the Pauling electronegativity scale, silicon has an electronegativity of 1.90 and oxygen 3.44. The electronegativity difference between the elements is therefore 1.54. Because of this moderately large difference in electronegativities, the Si\sO bond is polar but not fully ionic. Carbon has an electronegativity of 2.55 so carbon–oxygen bonds have an electronegativity difference of 0.89 and are less polar than silicon-oxygen bonds.

It is polar covalent, so has characteristics of both covalent and ionic bonds.

Bond order
Silicon–oxygen single bonds are longer but usually stronger than carbon–oxygen bonds, with a typical Si–O single bond length around 1.6 Å compared to 1.4 Å for C–O and with a bond energy of about 450 kJ/mol compared to about 360 kJ/mol for C–O. However, silicon–oxygen double bonds are weaker than carbon–oxygen double bonds due to a better overlap of p orbitals forming a stronger pi bond in the latter, with bond energies of 715 and 590 kJ mol−1, respectively. For these reasons, carbon dioxide is a molecular gas containing two C=O double bonds per carbon atom whereas silicon dioxide is a polymeric solid containing four Si–O single bonds per silicon atom.

Silicon-oxygen single bonds are by far the most common type of silicon-oxygen bond (see double bond rule), but silicon-oxygen double bonds can form in special circumstances. Silicon-oxygen double bonds are found in silanones. These compounds are normally very reactive and unstable with respect to oligomerization to siloxanes. They can be stabilised by coordination to a metal centre and by steric shielding.

Discuss true silanones, R2Si=O. Refs: H2Si=O (MeO)2Si=O, Ph2Si=O, Me(H)Si=O and Me2Si=O.

Silylated carboxonium ions, R2C=O+–SiR3, have been observed. These are a kind of silylium ion.

Bond strength
Bond strength vs. C–O, Ge–O, etc.? G&E p. 338: C–O is ~360 kJ/mol, Si–O is 452 kJ/mol.

Bond length
Does the length of the silicon-oxygen bond vary with the hybridisation of the oxygen atom? Yes - sp2 hybridised oxygen is found in Si=O bonds. Can you have a silyl oxonium ion like [H2C=O-SiMe3]+? If so, what is the Si=O bond length? One stable silanone has an Si=O bond length of 1.526 Å.

G&E p. 342: Si–O in alpha-quartz is 1.597 and 1.617 Å. In alpha-cristobalite, it's 1.61 Å. In vitreous silica, it's 1.58-1.62 Å.

March p. 25: sp3 C–O is 1.43 Å, sp2 C–O is 1.34 Å.

Bond angles
Si–O–Si notable for diverging from the typical C–O–C angle of 107–109°. Exception: Ph–O–Ph in diphenyl ether is (124±5)°. (ref - March for carbon)

G&E notes alpha-quartz Si–O–Si is 144°, beta-quartz is 155°, alpha-cristobalite is 147°, vitreous silica is (153±20)°, {Si2O76−} in thortveitite, Sc2Si2O7, is 180°. In Ln2Si2O7, it increases progressively from 133° to 180° as the size and CN of Ln decreases from Nd to Lu. In hemimorphite it is 150°. In lithium metasilicate and sodium metasilicate it is 134°.

G&E p. 613 summarises coordination geometry at oxygen. Linear Si–O–Si species include some silicates such as [O3Si–O–SiO3]6− in Sc2Si2O7, coesite (SiO2), and [O(SiPh3)2]. Bent Si–O–Si species are the norm, though.

Main explanations for larger E–O–E angle for Si vs. C:
 * Hyperconjugation (see below, p(O) → &sigma;*(Si–R) and to an extent p(O) → d(Si) both contribute)
 * Ionic effects (two partially positive Si atoms electrostatically repelling each other)


 * Oxygen difluoride F–O–F 103°
 * Dichlorine monoxide Cl–O–Cl 111°
 * Dibromine monoxide Br–O–Br 112°
 * Diiodine oxide I–O–I ???

Reactivity
Disiloxanes are less basic than ethers, i.e. worse donors of oxygen lone pairs. Accounts for physical and chemical properties of PDMS compared to polyethers.

Theory
From EJIC review 2021:
 * Older papers often invoked d-p pi backbonding but not (much) found in more recent calcs (silicon 3d orbitals too high in energy to accept oxygen lone pair electrons - at least unlikely to occur to a significant enough extent to dominate structure and reactivity)
 * Current main explanation is negative hyperconjugation (sigma backbonding): oxygen p orbital → Si–R &sigma;* (R = H, alkyl, aryl, etc.)
 * Effect is also present in ethers but stronger in siloxanes
 * Accounts for the high stability of polysiloxanes
 * Bent's rule (E–O–E oxygen lone pair has more s character so less basic when E = Si than when E = C?)
 * Can get an anomeric effect in siloxanes with more than one SiR2O unit, seen through conformational changes: O lone pair → Si–O &sigma;*
 * Polar covalent nature of the Si–O bond strengthens pi-backbonding from O to Si

Examples of compounds containing silicon–oxygen bonds

 * Silicon dioxide
 * Silicates / Silicate minerals
 * Silicon monoxide
 * Silicic acid
 * Orthosilicic acid
 * Silicotungstic acid
 * Siloxane
 * Silicones
 * Polydimethylsiloxane (PDMS)
 * Polymethylhydrosiloxane (PMHS)
 * Cyclosiloxanes
 * Hexamethylcyclotrisiloxane (D3)
 * Octamethylcyclotetrasiloxane (D4)
 * Decamethylcyclopentasiloxane (D5)
 * Drometrizole trisiloxane
 * Quadrosilan
 * Tetrakis(trimethylsilyloxy)silane
 * Sodium metasilicate
 * Silyl ether
 * Silyl enol ether
 * Organosilanols / Silanols
 * Trimethylsilanol
 * Diphenylsilanediol
 * Tris(tert-butoxy)silanethiol
 * Phenylsilatrane
 * Bis(triethoxysilylpropyl)tetrasulfide
 * Silanone
 * Silicon alkoxides
 * Tetramethyl orthosilicate
 * Tetraethyl orthosilicate
 * Silatranes
 * Silanones, which contain an Si=O double bond