Compilation: Bond vs. intermolecular interaction; chemical potential energy
Date: Wed, 26 Apr 2006
From: Andy Edington, Homestead High School, Wisconsin
Subject: Bond vs intermolecular force
Aren't both chemical bonds and chemical intermolecular forces just a balance between
electrostatic attractions and electrostatic repulsions? Wouldn't these both be considered chemical
potential energy?
-------------------Date: Thu, 27 Apr 2006
From: Guy Ashkenazi [a chemistry education researcher and Modeling Instruction consultant]
This is my point exactly. From my experience, the differentiation between chemical
processes and physical processes is artificial. It only works well for molecular substances: when
you evaporate a metal, is this a chemical or a physical process? Even there it has its problems:
when you evaporate water, are you just overcoming intermolecular interactions, or are you
breaking [hydrogen] bonds? Also, in physical chemistry, all changes in matter are treated as
chemical changes. For example, to calculate the vapor pressure of water, you basically calculate
Keq for the reaction H2O(l) <=> H2O(g). The only reason to make this distinction is
differentiating between evaporation and decomposition. Can anyone think of other points in
favor of this distinction?
One thing I would change, however, in Andy's statement is the phrase "electrostatic
repulsions". While the attractions in chemistry are all electrostatic, the repulsions are partly
electrostatic and partly quantum mechanical in nature. If it was just electrostatics, we wouldn't
need quantum mechanics to describe molecular structure!
Larry, can you post my paper on covalent bonding and the uncertainty principle to the
password protected part of the modeling web site, so people who do not have access to the
Chemical Educator could read it?
-------------------Date: Thu, 27 Apr 2006
From: Larry Dukerich ldukerich@MAC.COM, Dobson High School, Arizona State University
While the origin of these attractions are electrostatic in nature, it is useful to make a
distinction between them. Attractions are non-directional and non-specific. We generally use
this term to describe the interactions between particles when we are talking about phase changes.
Bonds are both directional and specific. We use this term to describe interactions during
chemical change. Usually (although not always) particles that are bonded require more energy to
separate than particles that feel attractions. I advocate making the distinction because beginning
chemistry students frequently have difficulty differentiating processes that involve separating
molecules from those that involve separating the atoms within molecules.
-------------------Date: Thu, 27 Apr 2006
From: Brenda Royce, University High School, Fresno, CA
I think you are right that both could be considered chemical potential energy with some
interactions falling in the 'gap' between them needing a more subtle distinction than we initially
use. The rationale for making a distinction between them is not that they are two very different
categories, but that there are different observable changes associated with each: intermolecular
forces are definitely involved when there are changes of state that do not include a change of
substance, while electron arrangements (or 'bonds') are involved with changes of substance
(chemical reactions). This is a distinction that is useful to beginning students who are forming
the conceptual structure for recognizing and describing changes in matter and energy transfer.
Whether or not the Ech and Ei classifications are used, I have found it to be very helpful to
give students a clear way of distinguishing the changes associated with these ways of interacting
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until they can discuss basic matter and energy changes well. On that foundation some of the
subtler distinctions can be discussed (H-bonding, types of intermolecular forces, enthalpy, and
such).
-------------------Date: Fri, 28 Apr 2006
From: Brant Hinrichs, Dept. of Physics, Drury University, Springfield, Missouri
I thought this was a pretty interesting comment. I was surprised by the description of
attractions as NON-directional and NON-specific. Is that primarily because attractions are
AVERAGED over all the different molecules interacting with any one particular molecule? So,
since there are molecules all around me, in general the net force they exert averages out? But
then how is there still attraction if it's an attraction. So I'm confused about how something can
be BOTH attractive AND non-directional? Whereas bonds are between two (or more) specific
atoms interacting to form that specific molecule?
--------------------Date: Sat, 29 Apr 2006
From: Guy Ashkenazi
I totally agree that bonds and attractions should be differentiated; I just wonder if doing
so by calling one "physical" and one "chemical" is the way to go. Maybe we should take an
energy approach here, and say that there are many mechanisms of attraction between particles
(covalent, ionic, metallic, H-bonding, Van der Waals), all of them electrostatic in principle, but
differ in detail, and therefore differ in energy. Some attractions are harder to overcome - require
more energy - while others require less. Therefore, when we have several of these acting together
in the same material, a small increase in average thermal energy (temperature) would cause only
the weaker attractions to break, where a much higher increase can ultimately cause all of them to
break.
Non-directionality means that Van der Waals interactions are attractive regardless of
molecular orientation. Two hydrogen molecules will attract one another when they are head-tohead, side-by-side, head-to-side, and whatever other angle. Non-specific means that a hydrogen
molecule will attract another hydrogen molecule, a chlorine molecule, a helium atom, and
whatever which has a positive nucleus surrounded by electrons. However, when hydrogen atoms
form bonds with an atom or a molecule, they will do so only with specific atoms (never with
helium, for example) and at specific proportions and orientations (v shaped with oxygen,
tetrahedral with carbon).
It is interesting to notice that according to this definition ionic and metallic attractions fit
the interaction definition better than the bond definition. Still we call them ionic bonding and
metallic bonding and not ionic interactions and metallic interactions. This only goes to show that
what really matters is the energy associated with the attraction, and not the details of the
mechanism.
-------------------Date: Sun, 30 Apr 2006
From: Larry Dukerich
The distinction between chemical and physical change is artificial. However, we choose
to impose structure on phenomena in order to help us better understand them. Since texts and
some standardized tests still ask students to make this distinction, I thought it worthwhile to
provide some simple structure and to consider the role of energy in these kinds of change.
Vanessa Barker's paper on misconceptions gave several examples of students' difficulties
in describing change:
An electric kettle was boiled in front of respondents so that bubbles could be seen
in the boiling water. They were asked "What are the bubbles made of?". The
replies included that the bubbles were made of heat, air, oxygen or hydrogen and
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steam. The question was answered by over 700 students and the same responses
were found. Proportionately, these varied from age 12 - 17 as follows:
heat
30% to 10%
air
30% to 20%
oxygen / hydrogen
25% to 40%
steam
15% to 30%
These data show that while the number offering a correct response, steam, does
increase between the ages of 12 and 17, most 17 year olds think either that water
can be split into its component elements by heating; or that heat is a substance in
its own right; or that air is contained in water.
Later she writes:
Students experience difficulty in recognizing when a chemical reaction occurs.
Many do not discriminate consistently between a chemical change and a change
of state, which chemists call a "physical change". Evidence for this comes from a
number of studies. For example, Ahtee and Varjola (1998) explored 13 - 20 year
olds' meanings for a textbook definition of 'chemical reaction'. Students were also
asked to state what kind of things would indicate a chemical reaction had
occurred. They found that around one-fifth of the 13 -14 year olds and 17-18 year
olds thought dissolving and change of state were chemical reactions. Only 14% of
the 137 university students in the study could explain what actually happened in a
chemical reaction.
Students are asked to try to imagine what is going on at the particle level during change
of all sorts. It seems to me that it is useful to distinguish between the generally reversible
changes in state which do NOT involve changing the identity of the particles from the often
difficult to reverse chemical changes which do result in new particles. This might be one of
those cases in which we provide a simple structure that helps the novice that we later need to
modify when we take a more careful look, when the student has a more robust conception of
what is taking place. After all, models do evolve as we examine phenomena at deeper levels.
What we have to decide is - are we doing the students any harm if, in their first pass through
trying to figure out what is going on at the particle level, we make this distinction.
-------------------Date: Mon, 1 May 2006
From: Andy Edington,Homestead High School
Larry's use of the word particles brought this to mind: I know what Larry meant, but
with three of my own children in elementary school, I can assure you that the use of the word
particles in this context is confusing to students (and many elementary teachers). Are particles
the little lumps (molecules) or the building blocks (atoms)? In my experience, the word particles
invokes in the students an image of atoms, much as the word legos invokes an image of the
individual blocks. This is also the case for the elementary school textbook definitions that are
drilled into my own childrens' heads. Consequently, if used in the stated form, Larry's definition
obscures rather than reveals meaning. To further confuse things, it is now becoming common
practice in chemistry education to talk about molecules of Fe(s) or NaCl(s).
One of my earliest teaching revelations occurred as I watched a veteran chemistry teacher
spend some time every day for weeks holding up plastic containers filled with homemade,
painted styrofoam ball models and ask compound or elemental, atoms or molecules, and pure or
impure? Concepts first, definitions last. Work with containers filled with lego models or ball
models to help students understand the distinction between a chemical change and a physical
change.
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Take a container of models representing water molecules and throw the water onto a
student. Have a student pick up the models and condense them back into the container. Have
yet another student chemically change the models into hydrogen molecule models and oxygen
molecule models. Have a group of students rapidly ignite the hydrogen molecule and oxygen
molecule mixture, changing them back into water molecules. While they are doing that, throw
another container of water on the students to stop the reaction.
If I break a stick in half, is that a chemical change or physical change? Is it reversible? I
would say that macroscopically, breaking a stick is a physical change, but microscopically;
breaking a stick is a chemical change. Physical vs. chemical change is a perfect chance to
compare and contrast and help students to start to see that, for the purposes of communication
and organization, man has put artificial distinctions on the natural continuum.
-------------------Date: Sun, 30 Apr 2006
From: Larry Dukerich
Guy asked me to post his paper (The Uncertainty Principle and Covalent Bonding) on the
modeling chem web site. You can find the link to it in the resources section just beneath unit 7.
-------------------Date: Mon, 1 May 2006
From: bRant hinrichs, Dept. of Physics, Drury University
Excellent, thanks so much for helping clear that up for me. So it is terminology in this
case that I got confused by. I like the word orientation much better than direction.
So if you'll allow me to test my understanding by pushing this a bit further, is it then safe
to say that this non-directionality occurs because of the underlying symmetry of the objects
involved - that is, because both hydrogens here have infinite rotational symmetry (I can rotate
them about any axis stuck through their exact center, and I can rotate as much or as little as I
like; but it still looks exactly like the same object - I can't tell I've rotated it just by looking and
comparing before rotation and after rotation), they have no preferred orientation in their
interaction.
-------------------Date: Mon, 1 May 2006
From: Guy Ashkenazi
A hydrogen molecule is not infinitely symmetric: it has an internuclear axis, which
defines a direction in space. There are non-directional and non-specific interactions BETWEEN
molecules. Within each molecule, the covalent bond between two hydrogen atoms is directional
and specific: the attraction of a hydrogen atom to the atom to which it is bonded is much stronger
than its interaction with any other atom.
Between two molecules, the interactions are called London dispersion forces, which arise
from induced asymmetry in the electron clouds of both molecules when they are close together.
This asymmetry is not specific, any molecule has an electron cloud that interacts in a similar
way, and the induced asymmetry can be induced by every other molecule in all orientations of
the molecule.
-------------------Date: Mon, 1 May 2006
From: Jessica Mamais (Ohio 2004), Chemistry Teacher, Olentangy High school
While we are discussing attractions and energy components, I need some clarification on
Hydrogen bonds. From my understanding a H-bond (attraction) is between a hydrogen on one
molecule and N, O, or F on another molecule. Yet, when explaining the difference in solubility
of NH3 and PH3, I have an answer to a text book saying that PH3 does not form H-bonds. Can
someone please explain the difference in the text book and my understanding?
-------------------Date: Tue, 2 May 2006
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From:
Larry Dukerich
As I understand it, a hydrogen atom bonded to a highly electronegative N, O, or F atom
acquires a partial positive charge due to the unequal sharing of the electrons in the covalent
bond. As a result it feels an attraction to the lone pair of electrons on a neighboring N, O, or F
that is considerably stronger than the normal London attraction or even a dipole-dipole
attractions.
According to Guy's definition (with which I agree), it is rightfully called a bond because
it is both directional and specific. The electronegativity difference between P and H is not so
large as to greatly skew the distribution of electrons in the molecule, so the partial positive
charge on the H atom is not nearly so great. Furthermore, the orbital containing the lone pair of
electrons in elements in the 3rd period or higher is more diffuse than is the case in N, O, or F, so
the attraction between the H and such "spread out" lone pairs would be weaker.
-------------------Date: Wed, 3 May 2006
From: Consuelo Rogers
It is easier to simply say that H-bonding is between molecules where a H-atom is bonded
with a SMALL, highly electronegative atom (i.e. N, O, or F).
-------------------Date: Thu, 4 May 2006
From: Guy Ashkenazi
It is more modeling-like to say that H-bonding is a term used to describe anomalies in
intermolecular bonding, such as high boiling points and good solubility in water, which cannot
be predicted based on simple periodicity or dipole moment values. These anomalies are typically
accompanied by a large change in vibrational frequency of one of the hydrogens in the molecule,
in most cases a hydrogen connected to a N, O, or F atom. We model this hydrogen as bridging
between two highly electronegative atoms, and attribute a partly covalent character to the bond
because it is specific and directional (~180o). There are exceptions to the NOF rule, such as the
poor water solubility of carbfuorides, and evidence of change in vibrational energy of HCl
interacting with ice (see http://www.nature.com/nature/journal/v417/n6886/full/417230a.html).
Always remember that it is just a model, and not a reason!
-------------------Date: Thu, 4 May 2006
From: Andy Edington,Homestead H.S
Here is a physical model to help students better understand hydrogen-bonding and the
extreme behavior we call Lowry-Bronsted acid-base behavior. Make some models of water
molecules using four styrofoam balls, two to represent the two lone-pairs and two to represent
the two shared-pairs (I use 4-inch diameter balls). Use three bulletin board push-pins, one to
represent the oxygen nucleus and two to represent the two hydrogen nuclei.
Holding two such water models up near each other makes it obvious why we would have
the super-strong attraction called hydrogen bonding. For acids and bases, an H+ ion is just a
push-pin. Students can see how it can easily pop on and off a water molecule. Using models like
these is called the tangent sphere approach.
If you like this, try making some models of oxygen molecules (O2) and nitrogen
molecules (N2). Again, use the styrofoam balls for the lone pairs and the shared pairs and pushpins for the nuclei. For oxygen, use one ball to represent the two shared pairs; for nitrogen use
one ball to represent the three shared pairs. Try making HCN. Students can quickly see how
double and triple bonds affect geometry.
I first saw this at the Woodrow Wilson Institute in Chemistry in 1987. The tangent sphere
approach was being used by Jerry Bell at Simmons College to teach organic chemistry in a
reasoning way (instead of memorization) by identifying nucleophilic sites and electrophilic sites
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and then predicting possible products. I know that at least one article about it has been published
in J.Chem.Ed. I believe it came out of the chemical bond approach (CBA) curriculum about the
same time as the CHEM-Study curriculum was being developed.
These models are great for comparing and contrasting with standard Ball and stick
models and space-filling models. Such comparisons really help students to become flexible with
their visualization skills. These models are also great for assessing if a student really
understands what the balls and sticks represent in a ball and stick model.
-------------------Date: Fri, 5 May 2006
From: bRant hinrichs
Is it possible to see this change in vibrational frequency in some kind of spectrum? Is it a
strong enough signal? Is it easy to measure/detect/differentiate (with NMR or some such)?
-------------------Date: Sat, 6 May 2006
From: Guy Ashkenazi
Of course. For example, in IR spectrum of alcohols: "The characteristic bands observed
for alcohols result from O-H stretching in addition to C-O stretching. The carbon-oxygen
stretching vibration of alcohols appears in a region complicated by many other absorptions, the
fingerprint region. The presence of a hydroxyl is better established by the O-H stretching.
The shape and frequency of an O-H band depends on hydrogen bonding. As hydrogen
bonding becomes stronger, O-H stretches appear at lower frequencies. In the vapor phase or in
dilute solution in non-polar solvents “free” hydroxyl group of alcohols absorbs strongly around
3600 cm-1. As the concentration of the solution increases intermolecular hydrogen bonding
increases and we see additional bands start to appear at lower frequencies, 3550-3200 cm-1 and
the “free” hydroxyl band decreases.
When the hydrogen of a hydroxyl group is involved in a hydrogen bond a resonance form
can be drawn in which the oxygen bears a negative charge. The contribution of this resonance
form reduces the single bond character of the hydroxyl bond shifting the absorption to a lower
frequency." (from http://www.yale.edu/ynhti/curriculum/units/1999/5/99.05.07.x.html)
The last sentence basically says that as the covalent character of the hydrogen bond
increases, the covalent character of the original O-H bond decreases, and a weaker bond has a
lower vibrational frequency. Same goes to NMR, where H-bonding shifts the resonance signal of
a proton to lower field (see
http://www.cem.msu.edu/~reusch/VirtTxtJml/Spectrpy/nmr/nmr1.htm).
-------------------Date: Fri, 5 May 2006
From: Consuelo Rogers
My students read about "shielding effect" as a reason for why Cl is not as electronegative
as F. They also read somewhere that these intermolecular forces (dispersion, dipole, and Hbonding) have an umbrella name of Van Der Waal's forces. What is all this terminology? What
is the history of these? What are the cut-offs for non-polar covalent, polar covalent, and ionic
bonds? I would prefer to describe bonding between atoms in ionic character.
-------------------Date: Sat, 6 May 2006
From: Guy Ashkenazi
The history traces back to Van der Waals equation of state for non-ideal gases, which
supposes some kind of attractive forces between gas molecules, which allows a description of the
liquification of gases. Since in 1873 Van der Waals knew nothing about electrons, he could not
identify the source of these attractions, so all intermolcular interactions which give rise to
liqufication were considered Van der Waals forces.
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Today it is customary to differentiate between hydrogen bonds and all other
intermolecular interactions: dipole-dipole, dipole-induced dipole, and induced dipole-induced
dipole (London- or dispersion-forces) interactions. My textbook refers to all three interactions as
Van der Waals intermolecular forces, and puts hydrogen bonding under a different section
heading.
Some textbooks, however, refer only to London forces (induced dipole-induced dipole
interactions) as VDW, and treat dipole-dipole as a separate class. I find no good reason for this
separation, because London forces are usually dominant even when the molecules have
permanent dipoles (compare the boiling points of CH3Cl and CCl4).
% ionic character only refers to covalent bonds! It is a measure of the contribution of
different electronic configurations to the overall covalent wave-function. It does not mean in any
way that the bond is partially ionic, in the sense of having substance properties which resemble
ionic substances.
For example, HF has a 59% ionic character (according to Pauling electronegativities), yet
it is completely molecular - it is a gas at room temperature, it does not conduct electricity in the
liquid state, it is a weak electrolyte in aqueous solutions, and in the crystalline state each
hydrogen has a single close fluorine neighbor, and each fluorine a single close hydrogen
neighbor (diatomic molecule). KI, on the other hand, has only 51% ionic character, yet it is
completely ionic: it is a solid at room temperature, it conducts electricity in the liquid phase, it is
a strong electrolyte, in the crystalline state each potassium has six close iodide neighbors, and
each iodide has six close potassium neighbors (NaCl lattice structure). The interpretation of the
low percent ionic character in this case is the expectation of high contribution of electron
sharing in a gas phase KI molecule.
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