The kind of questions that require the consideration of intramolecular attractive forces require comparison of boiling points or melting points...phase changes.

Phase changes in elements and compounds.

Elements

There are four type of elements;

Atomic (He, Ne, Ar, Kr);

molecular (H2, N2, O2, F2, Cl2, Br2, I2, P4, S8);

metallic (all metals);

extended covalent (C, Si);

For for atomic and molecular elements the bonding involved in phase changes are dispersion forces. Since the atomic and molecular elements exist as individual atoms or molecules the attractive force that must be overcome are intermolecular/interatomic. For metallic elements the forces that must be overcome are metallic. Covalent bonds must be broken in carbon and silicon. For carbon and silicon the issue of melting does require breaking covalent bonds, an intra-macromolecular force.

Compounds

There are two types of compounds;

Ionic (formula includes a metallic element and a nonmetallic element*);

Covalent (formula includes nonmetallic elements*);

For ionic compounds the attractive forces that must be overcome in melting are ionic, the electrostatic attraction of ions. This force, between ions, is an intra ionic type of attractive force. In an ionic compound the pure substance is a 3-dimensional array of oppositely charged ions. For covalent compounds the attractive forces that must be overcome are intermolecular and are either hydrogen-bonding, dipole-dipole and dispersion forces. For covalent compounds the smallest form of the pure substance is a molecule. Separating covalent molecules involves breaking intermolecular attractive forces.

*There are some exceptions to these general rules for using formulas to predict the type of bonding that occurs in the compound.

The most important exception to the rule that ionic compounds in their formula have a metallic element and a nonmetallic element, is the ammonium ion, NH4+. The presence of the pattern NH4 in a formula is also a characteristic of an ionic compound. Additional examples include other derivatives of ammonia (weak bases), including, methyl ammounium ion (CH3NH3+), dimethyl ammonium ((CH3)2NH2+), and others.

There are some exceptions to the formula of covalent compounds. There are examples of substances with formulas of metals and nonmetals which are principly covalent. This fundamental is related to the size of the cation. Most very small cations found in compounds show more covalent compounding. Examples are BeCl2, and Al2Br6 as well as other halides of these two metals. Magnesium show some covalent character also, but not to the same extend that beryllium does.

SiO2 is the only example of an extended covalent interaction in a compound. This substance has one of the highest melting points.

The distinction between intra- and inter- molecular attractive forces is important. Student must know that for covalent compounds, phase changes occur when intermolecular attractive forces are broken. Phase changes do not occur when intramolecular attractive forces are broken. Students must be sure to clearly state, or imply that it the attraction between molecules that are overcome when a phase change occurs.

Strengths of Intra- and intermolecular attractive forces

The strength of ionic bonds, which must overcome when ionic compounds melt, depend directly on the magnitude of the charge on the cation and anion, and inversely on the size of the ions. Charge is significantly more important than size when determining the stregth of ionic bonds.

The relative strength of intermolecular attractive forces in covalent compounds must be carefully compared. When considering compounds with elements in the first and second period only the order of strength is hydrogen-bonding > dipole-dipole > London dispersion forces. Remember dispersion forces are present in all substances. When we move to compounds containing elements from the third period or higher, we no longer have hydrogen-bonding occurring, only dipole-dipole forces (polar compounds) and London dispersion forces. In these compounds dispersion forces are generally the most important. So the more electrons in the compound, the more polarizable the compound is, the stronger its dispersion forces.

When comparing elements extend covalent attractive forces > metallic > dispersion forces. It is unlikely that students will ever have to argue which metal has a stronger attractive force. However, comparing extended covalent (C or Si) to dispersion forces (F2, Cl2, Br2) is likely. Extended covalent are alway stronger than dispersion forces! Just characterizing the presence of either and stating that extend covalent is stronger compared to dispersion is sufficient. The trend in strength of dispersion forces depends on the the number of electrons, the polarizability, of the element.