In chemistry, I guess we generally do not compare the strength of covalent bonds to that of ionic bonds.
However, take the case of AgI (silver iodide) vs PbI (lead iodide)
Now both have significant covalent character in their bonds as evidenced by the relatively small difference in electronegativities between the elements. But the Pb-I bond has higher covalent character and correspondingly lower ionic character.
From what I gather, PbI has a higher melting point than AgI. If we attribute this to the higher covalent character of the Pb-I bond, then covalent bonds are stronger than ionic bonds?
Melting does not break the bonds between the lead and the iodine within the molecule, it breaks down the intramolecular bonding. Wax has covalent bonds, yet a lower mp than the PbI or AgI.
Sorry, antechinus, but just about all of the info you have posted is incorrect.
Bond strength certainly correlates with melting point.
In ionic compounds, the ionic bonds are indeed broken in melting. One strong piece of evidence for this is the fact that molten ionic compounds (such as sodium chloride) conduct electricity, while the solid form does not.
Wax, a molecular compound, has covalent intramolecular bonds. These bonds are not broken when the wax is melted. Instead the intermolecular bonds are broken. These bonds are relatively weak, compared to the aforementioned bonds, and are termed London forces or Van der Walls forces.
A good example of a true covalent compound would be something like tungsten carbide, or solid carbon. Both have extremely high melting points.
And lastly, optimystique, your intuition is correct. Covalent bonds are indeed generally stronger than ionic bonds.
Well, alot of this has been touched already, but let me give my spiel along with an actual suggestion to help you figure out relative bond strength.
Melting point is mostly a result of intermolecular forces(dipole attractions,etc.), not intramolecular forces (bonds). When you melt something it doesn’t break the compound apart, it just makes it liquid. Melting ice doesn’t give you hydrogen and oxygen, it gives you liquid water. I suppose that when you’re melting salts it makes a sort of sea of ions, but alot of the bonding energy is still there in the form of electornic attraction. A better indication of bond strength would be looking at each compounds’ enthalpy of formation. Roughly speaking, that is the energy it takes for the compounds to form from their elements.
Robby, is there any reason why covalent bonds are generally stronger?
I thought that if you melt a salt, the energy required would be equivalent to the bond strength. In melted form the salt’s lattice structure would have broken down and you get a mixture of unbonded ions.
As for the enthalpy of formation, I couldn’t locate the data needed for the calculation e.g. first ionisation energy of silver, etc. so I can’t really comment on that. Anyone got the relevant data?
Intro chem concept: “compounds” fall apart easily, while “molecules” don’t. The former is ionic, the latter covalent, no? Are there ever any ionic bonds which form entities labeled “molecule?”
Covalent structures are much stronger than ionic compounds.
Mainly this is because of the nature of the bond, covalency comes about because of sharing an electron pair (sometimes 2, or even 3), which both atoms are attracted to. Whereas the only thing holding an ionic structure is the polarity of the atoms and magnetic attraction.
Diamond is a covalent macromolecular structure, hence its hardness.
Except that NaCl is only written as a molecule because it’s convenient. In reality, it’s just a bunch of Na+ and Cl- ions packed together in a 1:1 ratio. There is no actual “NaCl” molecule in there.
Introductory chemistry texts often categorize compounds into three categories: molecular compounds, ionic compounds, and covalent compounds.
The individual “units” in a molecular compound are molecules. The bonds within molecules (intramolecular forces) are covalent bonds, which are quite strong; the bonds between molecules (intermolecular forces) are relatively weak. (The latter forces include London forces, dipole-dipole bonding, and hydrogen bonds, depending on the compound in question.) Intermolecular forces for molecular compounds increase with the size of the molecule, leading to increasing melting/boiling points, for example, as you proceed from low molecular weight hydrocarbons to high molecular weight hydrocarbons. As mentioned, when you melt or boil the compound, you are not breaking apart the individual molecules, you are merely breaking the forces between molecules.
An ionic compound is made up of ions. There are no molecules to speak of. The formula written for ionic compounds simply describes the ratio of cations (positive ions) to anions (negative ions). The ionic bonds that bond the individual ions are much stronger than the intermolecular forces described in the previous paragraph. This leads to the high melting points observed for such compounds: e.g. 801 deg C for NaCl.
A covalent compound is made up of atoms. The very strong covalent bonds bond the entire structure together. To melt or boil the compound in question, these very strong bonds must be broken. This leads to very high melting points: e.g. ~3800 deg C for carbon.
One might think of the latter two compounds as being akin to “very big molecules” comprised of a great deal of individual ions/atoms. It is not very convenient to try to write a chemical formula showing this, however, as each sample would have a different formula! (For example, Na[sub]1,000,000,000[/sub]Cl[sub]1,000,000,000[/sub]) In practice, we simple write the ratio, NaCl.
The proper term for a single “molecule” [sic] of an ionic compound is a “formula unit.” For example, one formula unit of sodium chloride is comprised of a single sodium ion and an single chloride ion.
[sub]P.S. Bill, your electrical articles are fantastic. I used the concepts in them extensively when teaching physics two years ago.[/sub]
Paradigm, they are both of the same strength. Once a coordinate bond (where both electrons in the shared pair come from the same atom) is formed, it doesn’t matter whose electrons are whose.
oops, I didnt think I sent that post :smack: - realised too late you were only interested in ionic compounds.
When talking about bond strengths of ionic compounds it is best to think of lattice enthalpies. This takes into account the crystal structure of the compound, instead of thinking of a single ionic molecule in isolation. You can calculate the electrostatic interaction between the two ions using the charges and distance, but this is not really the bond strength for the substance in the solid state. It is OK if you are looking at a single molecule.
Yes there are ionic molecules, despite what some posters say. When a salt melts there is medium range order withinin the liquid. Molecules and clusters also exist in the vapour phase.
The ‘lattice enthalpy’ is the energy released when the gasseous ions combine to form the solid crystal. You are not comparing apples with apples when you compare ionic and covalent bond strengths, but this is the closest way of doing it, since the bond enthalpies of covalent bonds are also looking at a complete separation of the atoms. You do not get this separation when you simply melt the salt.
Looking at my SI data book, typical lattice enthalpies are around 1000 kJ/mol, whereas covalent bond enthalpies are around 200 - 600 kJ/mol. So looking at the question this way, ionic bonds are stronger.
The melting point PbI2 is less than that of AgI, but you are right to say that lead iodide has a more covalent nature.
Melting point of an ionic compound does correlate with the lattice enthalpy, however this involves the effects of the crystal lattice structure as well as the differences in electronegativities of the two ions. When you only look at the electonegativities the correlation is not there. For example, BeCl2 has a melting point of 405 degrees C. Compare that with PbI2, which has a similar melting point, but much closer electronegativities of the atoms.
Also compare GaCl3 which has similar difference in electronegativity to MgI2, but the former has a MP of 78 degC compared to the latter d700 degC.
You can calculate the lattice energy or molar internal energy fairly simply, taking into account the lattice type, but it is easier to look up the values in a table. btw according to the SI data book the first ionisation enthalpy of Ag is 737 kJ/mol.
ps Shimmery, some coordinate bonds are strengthened by pi-backbonding