Why is silicon dioxide acidic in nature and not amphoteric in nature?

Question

Silicon is a metalloid which means it has both metallic as well as non-metallic characteristics. But if silicon is a metalloid then why is silicon dioxide acidic in nature? If silicon has both metallic as well as non-metallic characteristics that means its oxide should have both acidic as well as basic characteristics which means that silicon dioxide should have both acidic and basic characteristics which means that silicon dioxide should be amphoteric in nature, right? Then why is silicon dioxide acidic in nature?

Answer 1

The question assumes that if silicon were a metal, it's $\ce{MO2}$ oxide would be basic. But in general $\ce{MO2}$ oxides of most metals are not purely basic but amphoteric. Titanium, for instance, is known to form titanates and its dioxide reacts with sodium hydroxide solutions even at ambient temperature [1]. Purely basic metal oxides almost always either have less oxygen than $\ce{MO2}$ or, where that stoichiometry is reached, are peroxides or superoxides (of $s$-block metals). Add the fact that silicon dioxide does show limited basic character through its reaction with hydrofluoric acid (WP does list the oxide as amphoteric), and we can fairly say that silicon dioxide does show intermediate characteristics between metal and nonmetallic oxides having the same stoichiometry.

Reference

1. Kostrikin, A.V., Spiridonov, F.M., Lin’ko, I.V. et al. "Interaction of components in the NaOH-TiO2 · H2O-H2O system at 25°C". Russ. J. Inorg. Chem. 56, 928–934 (2011). https://doi.org/10.1134/S0036023611060131

Answer 2

The acidic nature of $\ce{SiO2}$ comes from the $19$th century. In these old times, ions and electrons were unknown. Redox equations were not well established. The best known equations were the neutralizations corresponding to the general equation $$\ce{Acid + Base -> Salt + Water}$$ This equation is valid for all mineral and organic acids. And the "base", at this time, was what we call today our present hydroxyde or our metallic oxide. A base was a substance that turns the indicators basic, and reacts with an acid to produce a salt and water. An acid was a substance that turns the indicators acidic, and reacts with a base to produce a salt plus water. As a consequence, there was a tendency to generalize, and to include all inorganic substances in three categories : "acids", "bases", or "salts". Unfortunately, some substances like silica $\ce{SiO2}$ were difficult to put in any of these categories, as they do not react with acids and bases. But some day, somebody was able to make silica react with molten $\ce{NaOH}$ at high temperature, in an equation like $$\ce{SiO2 + 2 NaOH-> Na2SiO3 + H2O}$$ So here $\ce{SiO2}$ is not an acid, but it reacts with a base as if it was an acid : it produces a salt plus water after reaction with a base. This is why silica $\ce{SiO2}$ was classified as having an " acidic character".

Answer 3

SiO2 is acidic in nature due to a number of reasons:

1. Si has vacant d-orbital hence can act as Lewis acid(according to Lewis acid-base theory)
2. Si-O double bond is formed due to 2pπ-3pπ orbital overlap of O and Si respectively but due to the large size of Si, this Si-O double bond is not possible hence it completes it tetravalency by forming 4 Si-O single bonds(you will see that SiO2 I. E. Cassiterite is a 3D network solid in which SiO2 is in solid state and the Si here is sp3 hybridized). This Si-O single bond formed is much more stable because hybridized orbitals are much more stable than unhybridized orbitals due to lesser energy possesed by them hence Si hybrid orbital can overlap with p orbital of O This Si-O bond formed is more polar and covalent and any compound with polar and covalent bond(s) tends to be acidic This would be the explanation of acidity of SiO2 at the electronic level

Hope my arguments are reasonable enough to justify the question