This was one of my favourite types of questions when first learning the physics of chemistry. The answer is that there are several reasons as to why sodium chloride (NaCl) dissolves in water but not silicon dioxide.
- i) Sodium chloride
Let’s start by first thinking about sodium chloride (NaCl), more commonly known as salt, or table salt.
NaCl is an ionic compound. In fact, we can describe it in terms of its giant ionic structure. The ionic structure of NaCl is so big, we don’t exactly know how many ions there are. Such giant ionic compounds as in NaCl potentially consist of billions or even trillions of sodium ions and chloride ions compressed or compacted together. For this reason, some describe it as an infinite, or endlessly repeating, lattice of ions, such that the number of ions depends on the size of the crystal.
Furthermore, the giant ionic structure of NaCl can be considered as quite different from more commonly structured molecules, which contain an exact recipe of atoms. This is why NaCl is considered as a classic case of ionic bonding, in which atoms transfer or share valence electrons. This can be visualized using a Lewis Diagram:
Sodium chloride has a strong ionic compound. To invoke physics, we can say that this is because there is a strong electrostatic force between the oppositely charged ions (cations and annions). This is generally true when metals (in this case, sodium) react with non-metals (in this case, chloride).
We can explain these electrostatic forces, as well as the strength of sodium chlorides’ ionic compound, by citing Coulomb’s Law. This law of physics describes the strength of the force between two static electrically charged particles.
Despite its strong ionic compound, NaCl dissolves in H2O. The reaction looks like this:
NaCl(solid) + H2O —> Na+(aqueous) + Cl–(aqueous) + H2O
Notice that there is disassociation. The reason the sodium and chloride disassociate has to do with a number of factors:
Firstly, water is a polar molecule. In fact, it is very polar. This means that a water molecule has an asymmetrical arrangement of partial positive and partial negative charges that form polar bonds (below is a diagram that I have sketched for illustrative purposes).
As we know, water is made up of two hydrogen atoms and one oxygen atom. The atomic bonds in H2O are covalent bonds, which means that the electrons are shared. This results in the electrons that stay closer to the oxygen atom receiving a negative charge, while the hydrogen atom tends to have a positive charge.
NaCl, on the other hand, is made up of positive sodium ions and negative chloride ions. Hence, the polar ends of the water molecule attract their opposite charge parts of NaCl. More concisely, the positively charged water molecules attracts the negative chloride ions and the negatively charged water molecules attracts the positive sodium ions.
The reason salt dissolves in water is therefore due to how, the positively charged sodium ions are attracted to the negative polar area of the water molecule. Similarly, the negatively charged chloride ions are attracted to the positive polar area of the water molecule. These attractive forces with the water molecule overwhelm the forces between the positive sodium ions and the negative chloride ions, thus disassociation occurs and the ionic compound of NaCl goes into solution.
- ii) Silicon dioxide
But what about silicon dioxide (SiO2)?
Silicon dioxide has a giant covalent structure. This covalent structure, or macromolecules, is comprised of oxygen and silicon atoms.
The compound composed of silicon and oxygen to form SiO2 has a ratio of two oxygen atoms for every silicon atom. More concisely, each silicon atom covalently bonds to four oxygen atoms, while each oxygen atom is covalently bonds to two silicon atoms. In general, covalent bonds form when the element shares its four valence electrons, ns2np2, resulting in the formation of four covalent bonds (Clugston and Flemming, 2015).
Silicon dioxide, or silica, is very hard. Hence its diamond structure. This has to do with the strength of the covalent bonds, with oxygen atoms between each pair of silicon atoms. This strength depends largely on the electronegativity of the atoms insofar that electronegativity is the force between the electrons shared in the covalent bonding between the silicon and oxygen atoms.
Furthermore, SiO2 is not a molecule. It is a network covalent atomic solid. This giant lattice of covalently bonded atoms can be illustrated to look something like a 3D covalent network:
In that SiO2 is a network covalent atomic solid (under normal conditions), this giant covalent structure has very strong covalent bonds. These bonds are diffuse or spread throughout the structure.
The reason that there is strong electronegativity has to do with the oxygen atoms, which provide a stronger attractive force on the electrons than the silicon atoms, acquiring a partial negative charge. The electrons are also tightly compact, which means that SiO2 is not conductive (unless at molten temperature). On the other hand, SiO2 has high lattice energy.
All of this plays into why silicon dioxide is not soluble in water. In relation to bond in particular, SiO2, or silica sand, is insoluble because the attractive forces of the water molecules are not strong enough to break the covalent bonds between the silicon and oxygen atoms. In more precise terms, there is no attraction between the polar water molecules and the silicon or oxygen atoms due to the non-polarity of SiO2. This is because, despite the silicon-oxygen bonds being very polar, the geometry of the molecule – there are four silicon-oxygen bonds that cancel the dipole – means the dipole moments cancel resulting in non-polarity.
In conclusion, while sodium chloride (NaCl) dissolves in water due to the attractive forces with the polar water molecules overwhelming the forces between the positive sodium ions and the negative chloride ions, resulting in disassociation; silicon dioxide (SiO2) does not dissolve due to being a giant covalent structure in which the dipole moments cancel resulting in non-polarity.
Atkins, P., and De Paula, J. (2013). Elements of Physics Chemistry. Oxford University Press. Oxford, UK.
Clayden, J. Greeves, N., Warren, S. (2012). Organic Chemistry. Oxford University Press. Oxford, UK.
Clugston, M., and Flemming, R. (2015). Advanced Chemistry. Oxford University Press. Oxford, UK.
Lister, T. and Renshaw, J.. (2015). AQA Chemistry. Oxford University Press. Oxford, UK.
Weller, M., Overton, T., Rourke, J., and Armstrong, F. (2014). Inorganic Chemistry. Oxford University Press. Oxford, UK.