Model kits with two red oxygen atoms, two green fluorine (or chlorine atoms) atoms, two blue nitrogen atoms, two black carbon atoms, 8 white hydrogen atoms and 3 different strength rubber bands.  Students build models, sketch diagrams, identify the location of IMFs between molecules, estimate boiling points.

Author: T. Greenbowe   This page is under construction.

Part 1   ​
Two molecular models of each of the following compounds are built: 
 CH₃CH₂NH2 , CH₃CH₂F , CH₃CH₂OH .

"Which liquid sample of the above three compounds has the higher boiling point?"  Explain.

Three different types of rubber bands (different strengths) are used  to represent different strengths of IMFs.  The rubber bands are used to connect two molecular models together at the appropriate partial negative charged atom in one model - - -IMF - - - and the partial "+" charged atom in the other molecule. Students draw Lewis Structures of the two models near each other and use a dotted line to indicate the location of the IMF.

a.  Each group builds one CH₃CH₂NH₂ model.

b. Use your Table of Electronegativity Values of the Elements to determine which atom(s) has/have a partial positive charge (d+) and which atom(s) has/have a partial negative charge (d+).  Determine the DEN between the atoms having the d+ charge and d- charge.

c. Bring the two models closed together.  You have three different rubber bands, each representing a different strength IMF.  Use the correct rubber band to connect the two CH₃CH₂NH₂ models by the primary IMF at the correct atom locations.

d. How many possible different arrangements are there?

c. Identify the IMF.  Classify the strength of the IMF as weak, medium, or strong.

d. Sketch two 3D Lewis structures representing the two models and use a dashed line to show the correct location of the IMF.

Physical Properties Data Tables

Part 2  ​
Two molecular models of each of the following compounds are built: 
 CH_4 , CH₃Cl , CH₂Cl₂ , CH₃OH , CH₂(OH)₂

"Rank order the strengths of the IMF found in the liquid state of each of the above compounds.  Rank order each liquid sample of each compound with respect to increasing boiling points."

a variety of rubber bands in different colors and different strengths to represent different strengths of IMFs.  The rubber bands are used to connect two molecular models together at the appropriate partial negative charged atom in one model - - -IMF - - - and the partial "+" charged atom in the other molecule. Students draw Lewis Structures of the two models near each other and use a dotted line to indicate the location of the IMF.

Student Data Table

Data Table

copyright 2018 Chemistry Education Instructional Resources and the Department of Chemistry & Biochemistry, University of Oregon, Eugene, Oregon USA

A set of Power Point slides is available to accompany this activity.

Curriculum Notes 

 Student Difficulties (Misconceptions) Related to IMFs

1.  Students exhibit difficulty with the concept of an Intermolecular Force being a force of attraction, an electrostatic interaction, between two or more molecules.  The majority of students will identify an IMF as one of the chemical bonds within a molecule (intramolecular force of attraction), or the identification is ambiguous or contradictory.  Students confuse the terms inter- and intra-molecular forces.

2.  Students have difficulty visualizing a three-dimensional molecule from a two-dimensional Lewis Structure.

3.  Students have difficulty locating atoms in a polar molecule that have a partial positive or partial negative charge.

4.  Students have difficulty drawing structures and showing the location of the intermolecular force of attraction between molecules.

5. Students do not give the right reason for a molecule having stronger LDFs in comparison to another molecule (e.g., I₂ has stronger LDFs than Br₂ because it has a higher molar mass [instead of discussing how the larger polarizable electron cloud in I₂ is able to generate more LDFs and stronger LDS than the smaller electron cloud in Br₂]).

6. Students often use only three IMFs found in samples of pure substances (solids and liquids) - Dipole-dipole, Hydrogen bonding, LDS -rather than the IMFs found in solutions (ion-dipole, ion-indiced dipole, dipole-indiced dipole, etc.).

7. Students make the mistakes of arguing that dipole-dipole forces are ALWAYS stronger than LDFs, and comparing LDFs based on mass of atoms rather than polarizability.

8. Students confuse “higher condensation point” with “lower boiling point” .

9. When asked to compare two pure liquids each in its own sealed container at the same temperature, students associate "low vapor pressure" with weak intermolecular forces.

Learning Objectives for this Activity

1.  Identify the atoms in a model of a molecule having a partial positive charge or a partial negative charge: bond polarity.

2.  Acknowledge the fact that intermolecular forces are electrostatic force of attraction (and repulsion) between two or more molecules.  The forces responsible for keeping molecules or atoms intact as a solid or liquid are  intermolecular attractive forces. 

3. Given a compound or element that is a pure solid or liquid, identify the types of intermolecular attractive forces that occur between two or more molecules or atoms.  IMFs in pure substances: Dipole-dipole, Hydrogen Bonding, London Dispersion Forces.

4. Identify the location(s) and type of an intermolecular force between two or more molecules as represented by two molecular models.  A rubber band connects a partial positive charge atom in one molecule to a partial negative charge atom in another model.  \

5.  Using a Table of Electronegativity Values of Elements determine the relative strength of the IMF between molecules.  Different types of rubber-bands are used to represent IMFs of different strengths. 

6. Illustrate using Lewis structure diagrams or 3D projection diagrams of molecules, the intermolecular force between two or more molecules (or atoms).

7. Identify and correlate the types of IMFs and the strengths of IMFs present in a sample of liquid or solid with the physical properties of that compound or element: melting point, boiling point, heat of vaporization, etc.  The strength of the IMF(s) interactions between molecules (or atoms) in a solid or liquid determines its physical state and determines its physical properties.

8. Given three compounds, rank order the compounds in terms of increasing physical properties such as boiling point, melting point, or vapor pressure.

9. Students have the opportunity to construct and revise representations, models, and explanations that allow them to predict and explain phenomena.

AP Chemistry Enduring Understanding 5D

Electrostatic forces exist between molecules as well as between atoms or ions, and breaking the resultant intermolecular interactions requires (an input of) energy.

AP Chemistry Learning Objectives

LO 2.16  The student is able to explain the properties (phase, vapor pressure, etc.) of small and large molecular compounds in terms of the strengths and types of intermolecular forces.

LO 5.9 The student is able to make claims and/or predictions regarding relative magnitude of the forces acting within a collection of interacting molecules based on the distribution of electrons within the molecules and the types of intermolecular forces through which the molecules interact.

Literature Articles Related to Student Difficulties with IMF

Copper, M; Williams, L. C., Underwood, S. M. Student understanding of Intermolecular Forces: A multi-modal study.  Journal of Chemical Education, v92 n8 p1288-1298 Aug 2015. (EJ1073053)

Quotes:  "It is our contention that, to develop a robust understanding, the curriculum must be restructured to emphasize the connections between important ideas and that students must be given opportunities to reflect on and make their thinking visible (Ref#45)  That is, students must have the opportunity to construct and revise representations, models, and explanations that allow them to predict and explain phenomena. Otherwise, it becomes too easy to assume that students have learned important concepts because they can choose the correct answer on an examination."

 "If students are not ever asked to write and draw, to reflect, to explain, and to revise their ideas, but instead are only assessed by which item they choose on a test or randomly generated homework, it is unlikely that they will develop a robust and coherent understanding of core concepts. This is not to say that multiple-choice items are never useful  but that students must also be given many opportunities to construct answers for themselves as they learn.."

Three days of lead time is required for this project.

Discussion 

Students work in groups to draw Lewis structures and to construct two 3D molecular models (ball-and-stick) of a given compound, for example, ethanol CH3CH2OH.

Students a. identify the partial positive charge atoms and the partial negative charge atoms in each model that are involved in the IMF, b. calculate the difference in electronegativity between the opposite charged atoms, c. determine the types of IMF involved in a liquid or solid sample. Students connect their two models by rubber bands representing the location of an intermolecular force.  For example, in ethanol students would connect the partial positive charge hydrogen atom of one molecular model to the partial negative charge oxygen atom in the second molecular model.

Students select the appropriate rubber band width to correlate with the strength of the IMF between models.  For example a thin rubber band is used to represent the IMF between two ethyl fluoride, CH3CH2F models while a thicker rubber band is used to represent the IMF between two ethanol models.  The major IMF holding ethanol molecules together is called hydrogen bonding and this is a stronger IMF compared to the dipole-dipole IMF in a liquid sample of ethyl fluoride.  Students correlate the strength of the IMF to a physical property of the liquid state of the compound -  boiling point, vapor pressure,  and Heat of Vaporization.  For example, ethanol has a higher boiling point compared to ethyl fluoride due to the stronger hydrogen bonding IMFs in a sample of ethanol.