Ending Misconceptions About Inter- and Intramolecular Forces

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Authored by

  • George Miller
    University of California
    Irvine, California

Understanding Physical and Chemical Changes

There is often confusion on several issues relating to inter- and intramolecular forces.

Physical change can involve changes in intramolecular forces (chemical bonds). The strength of the chemical bonds depends only on their type (ionic or covalent).

When asked to explain differences between various substances as they are heated, students can confuse physical and chemical changes in the substances. Further, students, given the opportunity, will incorrectly assign relative boiling points to substances. For example, they may assert that methane has a high boiling point because it contains strong covalent bonds. Alternatively, they may overgeneralize, such as arguing that diamond has a lower melting point than salt because ionic bonds are stronger than covalent bonds. They may incorrectly use mass arguments, such as that water has a higher boiling point than ammonia because water molecules (mass 18) have a higher mass than ammonia molecules (mass 17). While mass may be part of the story in terms of correlation (e.g., bromine is more massive than fluorine), the important factors always involve attractive forces.

Why do these misconceptions exist? This appears to be founded in a lack of understanding that more than one kind of "force," or interaction, can be occurring in one substance at the same time. Furthermore, there is a lack of appreciation of the relative magnitude of the forces within substances and modest levels of temperature increase as a result of heating. Some students appear to have insufficient practice in thinking of "two variables at once" and in calibrating their thinking with regard to energy and forces. Unfortunately, lack of familiarity may be responsible for students reporting that water splits into hydrogen and oxygen (something they may have seen once in an electrolysis demonstration) when heated to its boiling point.

Linguistic issues do not help, causing some students to reverse the terms. One dictionary says inter- can mean "between, or among" but also "within," whereas intra- means "within" or "during"! The word internal seems to relate to inter- when the appropriate term is intra-, and "intermolecular" forces are actually "external" to the molecules! Inter- in the latter case stands for "between," not "internal."

What is the correct picture? Attractive forces hold molecules together internally and with one another. In all substances—except the noble gases—strong, stable chemical bonds hold atoms together in assemblies called molecules. These require significant amounts of energy as well as suitable pathways, such as catalysts, to "break."

Energy often has to come from formation of stronger bonds, as when hydrogen gas burns in air to form water, or a surface catalyst such as nickel metal permits hydrogen molecules to break apart and hydrogenate an olefin.

Modest amounts of energy added to a solid substance will cause melting and eventually boiling of that substance. In these transformations, intermolecular forces arising from generally dynamic interactions between the moving molecules are given energy so that further motion and separation of the molecular units from one another is possible. The amounts of energy involved in the physical change are an order of magnitude smaller than those involved in breaking bonds (e.g., 40.7 kJ to vaporize 1 mole of water versus 934 kJ to break 1 mole of O-H bonds).

The various types of forces, both intermolecular and intramolecular, vary significantly in strength both within their category and between one another. It is thus a serious oversimplification to make general conclusions about material substances without detailed consideration of all the various force types that are present. Many texts treat this fairly well but often rather briefly. Even charts constructed to help students identify the types of forces are accompanied only by a statement of the general trend in strength, and covalent bonding may not be included in the list. The additional complication to be mastered is the formation by many substances of network solids involving multiple bonds of either ionic or covalent type. Rarely is a single covalent bond directly compared to a single ionic attraction in an ion pair, the only true comparison of "strength." This has to be done to clarify the diamond or sand comparison to salt, for example.

Texts usually list the forces in order of increasing strength as London dispersion (induced dipoles), dipole-dipole, hydrogen bonding, ion-dipole, ionic, and covalent. While this is a general trend, many exceptions exist, and only continued discussion will bring familiarity with when the general trend is sufficient as explanation and when more detailed consideration is needed.

Instructional Recommendations

Standard textbooks contain adequate and careful explanations of the correct picture. Students need to be made more familiar with energy transfer and substances in which it is frequently observed to cause physical changes. In addition, the possible existence of multiple forces within an assemblage of molecules must always be considered.

The latter can usefully be emphasized with the use of molecular models. However, models are usually used to detail "intra-" chemical bonding (ball and stick) rather than intermolecular interactions. Ideally the atoms in model kits should be equipped with weak Velcro to allow then to stick together weakly to illustrate the weaker forces. One possibility is to use Lego bricks firmly pegged together to form the molecules and then hold assemblies of the molecules together with loose elastic bands. The pegs represent the intramolecular covalent bonds, and the elastic bands represent the weaker intermolecular forces holding together a model of a molecular solid that will have low melting and boiling points compared to a fully interlocked covalent network solid. Emphasize that just pushing such models around (e.g., shaking them up in a Ziploc bag) is modeling adding heat, and it may "break" or loosen the weaker elastic bands but will not break the tight peg "bonds."

Students need practice in "explaining" these issues and considerable dialogue with teachers and one another about what is a satisfactory explanation. They have to be encouraged to take the explanation to the core of the issue rather than stay on the surface. One of my favorite examples is to ask students to explain why baseball team X won the game against team Y and to test their explanation against that of a baseball fanatic. The explanation that team X won because they scored more runs than team Y is perfectly accurate, but it is a seriously insufficient explanation for the baseball fanatic, who would want to know about each pitch and its result. They should assume their explanations for chemistry are to be read by an equivalent chemistry fanatic.

The AP Chemistry Examination often simplifies the question by presenting closely similar molecules for comparison of physical properties (such as NH3 and NF3 in 2005, F2 and I2 in 2004, and NH3 and CH4 in 2001). However, this may not always be the case, as in 2001, when students were asked to explain why Si melts at a much higher temperature (1,410°C) than Cl2 (-101°C).

While the scoring guidelines may be simple, to "get the points," the best students should be able to answer this somewhat as follows:


Si is a network solid with a diamondlike structure (though weaker than diamond), with each silicon atom bonded to four neighboring silicon atoms. When heated, a large number of moderately strong bonds must break for the atoms to be free to move and the solid to melt.

In contrast, each chlorine atom forms one strong covalent bond to a second chlorine atom to form a Cl2 molecule. The diatomic molecule formed is symmetrical and thus nonpolar. Weak induced-dipole forces (London dispersion forces) resulting from temporary distortions of the electron clouds on the molecules are the only attractive forces between one Cl2 molecule and the next. At low temperature, with very little heat added, these molecules are thus free to move and behave as a liquid. Hence the low melting point of Cl2.


Note that it is important to fully discuss both the intermolecular and intramolecular forces in both species that are being compared.

Dr. George E. Miller is senior lecturer SOE emeritus in the Department of Chemistry at the University of California, Irvine, where he is also the principal scientist and the reactor supervisor of UCI's nuclear reactor and faculty director for science education programs at UCI's Center for Educational Partnership, including Faculty Outreach Collaborations Uniting Scientists, Students and Schools (FOCUS). He was a member and chair of the AP Chemistry Development Committee and is currently a member of the College Board Subject Test in Chemistry Development Committee. He has been a Reader, Table Leader, and Question Leader for AP Chemistry.