Ionic Vs Molecular: Fill & Download for Free

I’m not sure if ionic compounds like NaCl boil at all without decomposition. would sodium chloride vapor consist of gaseous sodium and chloride ions, or would some chemical change need to occur? I don’t know. However, the difference in typical melting temperatures is easily explained by considering the kind of forces that hold ionic vs. molecular crystals together. Electrostatic forces between + and - charges are responsible for the structure and stability of ionic solids. Such forces are strong at the kinds of distances separating them inside a crystal. Molecular crystals are stabilized by relatively weak forces. Van derWaals forces and even weaker intermolecular forces are what cause covalent compounds to pack closely in molecular crystals. Furthermore, such forces decrease with the inverse 6th power and I believe, even the 12th power. Consequently, it takes less thermal energy to overcome these forces, compared with electrostatic forces in ionic solids.

Well, metal-nonmetal aren’t always ionic, but some of our favorite compounds have alkali metals which easily give up their electron, and halogens which really want one.But there are metal-nonmetal compounds that are considered covalent.Nonmetal+nonmetal is most often a molecular element, such as H2, O2, N2, Cl2, which are necessarily covalent, as the two are equal.Otherwise, ionic vs. covalent is more of a degree. The III-V compounds that are some of our favorite semiconductors are not very ionic.And hydrogen is a non-metal that forms ionic bonds in some cases.

An ionic bond is defined to be a bond that is sufficiently lopsided that one set of atoms has all of the electrons donated to them by another set of atoms. However, there is no such a thing as a purely ionic bond! Even crystals made of ions have partial covalent character in the bonds or interactions between atoms making up the crystals. Traditionally, an ionic bond appears when the difference between electronegativity of the interacting atoms is greater than 1.7 Paulings. Covalent bonds containing atoms whose electronegativities differ by 1.0 Paulings are considered to be polar covalent bonds. Thus, ionic bonds can and do have partial polar covalent character. In bonds between atoms with little or no electronegativity difference, the bond is considered to be covalent but non-polar.Many factors can change the ionic bond character of a bond. Moving atoms close together can increase the covalent bond character and change the hybridization of orbitals being used to form that bond. Placing the atoms farther apart will separate them as distinct ions or fragments, with greatly reduced covalent bond character. The environment around an atom can also affect the bond properties of atoms attached to it. Atoms dissolved in solvents are interacting with solvent molecules, and will have different bond character compared to those in the gas or plasma state.Electron density is a common way of recognizing possible ionic bonds, as ionic bonds tend to be omnidirectional (making them great for crystal structures) and have a spherical appearance. When we determine the electron density via computational chemistry or X-ray crystallography, we are looking at a specific concentration of electrons surrounding the atoms or molecule. This is called an isosurface, with the amount of electrons an isovalue. Isovalues are defined with the unit being electrons per cubic angstroms. Changing the isovalue will give us surfaces that look different for the same system. Lower isovalues represent low concentrations of electrons, and tend to be farther away from the atoms, while high isovalues describe high concentrations of electrons, being close to the atoms. If we were to use low isovalues on some molecules, crystals, and systems, they will look like covalent bond systems, even if they contain atoms with large electronegativity differences. To illustrate this point, we take GeF4, which contains germanium and fluorine. In the Pauling scale of electronegativity, Ge has a value of 2.01 Paulings, while fluorine has 3.98 Paulings. This should give us a difference of 1.87 Paulings, enough to create ionic bonds between these atoms.GeF4 is a tetrahedral molecule with Ge-F bonds measuring 1.644 angstroms long. It is the 'ionic' analogue of methane, CH4, and we find some surprises when we look at its electron density (calculated with density functional theory at the Becke-3-Lee-Yang-Parr level), below.This is what its electron density looks like, at an isovalue of 0.08 electrons per cubic angstrom. It almost looks ionic, but still has significant overlap between the individual atoms. We also get polar covalent bond appearances when we try lower isovalues for the electron density calculations, below.Medium electron density, at 0.01 electrons per cubic angstroms. It still looks quite like a typical polar covalent molecule, much more like CF4 or CCl4.Low electron density, at 0.002 electrons per cubic angstrom. Notice that decreasing the isovalue of electrons per cubic angstrom produces a bigger surface. Such surfaces give an appearance of covalent bonds, even in ionic materials or solids. Thus, GeF4 does not behave purely as an ionic compound, as its electron density looks much like those of polar covalent compounds.Sodium chloride is a different story when we look at its electron density. Sodium has an electronegativity value of 0.93 Paulings, while chlorine has a value of 3.16 Paulings. This gives us a difference of 2.23 Paulings, more than enough to fulfill the traditional definition for ionic bonding. We can easily examine the bonding in this compound by looking at a molecule of NaCl, representing two atoms out of billions or more in a typical salt crystal. The high isovalue electron density surface of a NaCl molecule is below.Notice the spherical shape of the electron density around these two atoms (Na is on the right side, Cl is on the left side)? That is a prominent feature of ionic bonding! The two atoms are 2.372 angstroms apart, giving them ample spacing to develop ionic bonding. However, if we decrease the isovalue used to obtain the electron density surface of the NaCl molecule, we begin to get the covalent bond appearance as the concentration of electrons decreases and we go farther away from the atoms. The medium and low isovalue surfaces for NaCl are shown below:Medium electron density, already showing features of polar covalent bonding. Na is on the left, and Cl is the right atom here.Low electron density surface, showing more of a polar covalent bond than the ionic bonding we saw earlier. It is possible that all ionic materials have a particular isovalue of electron density where the ionic bonding and covalent bonding begin to blur together. This isovalue produces surfaces with electron density from separate atoms barely touching each other, and it is very sensitive to distance, number of coordinating atoms, environmental effects, and even isotope substitution.Another issue with the 'pure ionic bond' is that some purely homonuclear bonds can actually look ionic! These bonds are made of identical atoms, with no electronegativity difference to polarize any of the atoms. Disodium, Na2, is such an example. In a plasma or gas of sodium, molecular forms of this metal can exist or survive, and it has a distance of 3.086 angstroms between the two sodium atoms. If we look at the high electron density surface for this molecule, we find spherical distribution around the atoms!Without knowing this is a homonuclear bond, we would have guessed this was an ionic bond. When we decrease the isovalue for electron density in disodium, we do find blurring of ionic and covalent bond behavior, just as we did with the molecular NaCl species, below.Here, the spherical shapes give way to a more covalent bond electron distribution. We see more of the covalent bonding when we further reduce the electron concentration being probed around the atoms.Thus, we cannot completely eliminate covalent bonding in any assemblage of atoms. It is going to be there, regardless of how lopsided the electrons are going to be shared among the atoms. The isovalue surfaces are a proof that no pure ionic bonds exist. An ionic bond is always partially covalent. This finding also applies to coordinate bonds, as those between boron and nitrogen in borane-amine complexes. The molecule, BH3NH3, is a good model to examine the B-N bonding, which is traditionally considered to be a dative bond. The nitrogen atom gives two electrons to the boron atom, and alters the formal charges among the boron and nitrogen atoms. When we take a look at the electron density of this complex, we find that the B-N bond is different compared to the B-H and N-H bonds, as the electronegativity difference between B (2.04 Paulings) and N (3.04 Paulings) is greater than those of B vs. H and N vs. H. The B-N bond has a difference of 1 Pauling, so it is considered to be a polar covalent bond.The dative bond has made the polar covalent bond look almost ionic in this high isovalue surface, but the short distance between the B and N atoms (1.842 angstroms) makes perfect spherical distributions of electrons difficult to achieve. Lower isovalues give us a better view of the polar covalent bond.At this isovalue, we are not able to easily distinguish BH3NH3 from ethane, C2H6, whose electron density surface is below. Ethane has a C-C bond that is 1.512 angstroms long, and is not considered to have strong polar bonds.With all things considered, the important thing to understand is that no single bond type dominates completely. A bond can be ionic, covalent, metallic, and polar all at once!