Experiment 04/05 Two experiments in connection with intramolecular and intermolecular forces By Christian Redeker 10.11.2007 Two experiments in connection with intramolecular and intermolecular forces 2007 Contents
1.) Hypothesis ............................................................................................................................3
2.) Diagram ................................................................................................................................8
3.) Method..................................................................................................................................9
3.1) The miscibility of liquids ................................................................................................9
3.1.1) Apparatus .................................................................................................................9
3.1.2) Saftey Notes .............................................................................................................9
3.1.3) Procedure................................................................................................................10
3.2) Testing liquids for a permanent dipole .........................................................................10
3.2.1) Apparatus ...............................................................................................................10
3.2.3) Saftey Notes ...........................................................................................................10
3.2.2) Procedure................................................................................................................10
4.) Errors ..................................................................................................................................11
4.1) The miscibility of liquids .............................................................................................11
4.2) Testing liquids for permanent dipole ............................................................................11
5.) Results ................................................................................................................................11
5.1) The miscibility of liquids ..............................................................................................11
5.2) Testing liquids for permanent dipole ............................................................................13
6.) Conclusion/Evaluation........................................................................................................14
6.1) The miscibility of liquids ..............................................................................................14
6.2) Testing liquids for a permanent dipole .........................................................................18
7.) Evaluation...........................................................................................................................20
8.) Bibliography .......................................................................................................................20
8.1) Notes .............................................................................................................................20
8.2) Figures...........................................................................................................................20
2 Two experiments in connection with intramolecular and intermolecular forces 2007 1.) Hypothesis
Intramolecular forces respectively bonding forces between atoms and intermolecular forces,
forces between molecules, are the foundation of the occurrence of any chemical reaction and
to the chemical and physical properties of substances. Both, intramolecular and
intermolecular forces have their origins in the attraction of charged particles, namely
negatively charged electrons and positively charged protons which together make up an atom.
Both principles of both types of forces, intramolecular and intermolecular, are explained in
the following paragraphs.
Bonding between atoms is the fundamental principle of chemistry. “Chemistry is the study of
the elements and the compounds formed when they bond with each other”1. However, the
elements of the periodic table do not bound with each other randomly. Whether an atom or a
molecule binds to another atom or molecule is influenced by many variables. However,
principally a chemical reaction occurs to lower the energy of the involved atoms or
compounds. Those reactions are said to be energetically favourable because they cause a
decrease in energy of the atoms. Every atom strives for having the lowest energy possible.
The energy an atom has depends on the arrangement and of the number of charged particles,
proton and electrons, which make up the atom. The lowest energy an atom can obtain is
principally reached when the valence shell of the atom has the highest possible number of
electrons. This connection is called the octet rule. It is rather a rule of thumb than an accurate
description, but it works in many cases.
It can therefore be stated, that atoms usually combine to obtain a full valence shell.
Principally, there are two ways how atoms can achieve a full valence shell through bonding.
Those are ionic bonding, in which electrons are virtually transferred from one atom to
another and covalent bonding, in which electrons are shared between atoms.
Ionic bonding occurs between two atoms which have a high electronegativity difference
(usually above 2). Electronegativity is measure of how strong an atom can pull electrons to
itself. The electronnegativities of elements from group 1 are very low while the
electronegativity of group 7 elements are very high. Groups in between those two groups
have steadily increasing values of electronegativity, so that group 1 and group 7 elements are
virtually the extreme sides in terms of electronegativity. Furthermore, the electronegativity
decreases down the groups, so that caesium has the lowest electronegativity (0.8) and fluorine
has the highest electronegativity (4.0). The realtion between the electronegativities of the
elements of the periodic table is visualised in figure 1.
Ionic bonding therefore only forms between atoms which are far apart in the period table. In
view of the elements that are rather on the left hand and side of the periodic table being
metals, and the elements on the right hand side being rather non-metals, it can be stated that
ionic bonds can only form between metals and non-metals. The exact definition whose
elements are metals, non-metals and metalloids are shown in figure 1.
3 Two experiments in connection with intramolecular and intermolecular forces 2007 Figure 1: The electronegativities for the elements of the periodic table
As stated above, ionic bonding involves electron transfer from one atom to another. As a
consequence of this electron transfer ions, particles which have an overall charge, are created.
The electron donor becomes thereby positively charged, and is therefore called a cation,
while the electron acceptor becomes negatively charged and
is therefore said to be an anion. The two charged particles
exert forces on each other due to their charge. Compounds
which are formed by an ionic bond arrange therefore when
solid in a lattice, in which “ions pack together in an array of
alternating opposite charges”2. The principle arrangement
of an ionic lattice is shown in figure 2. As a consequence of
the electron donation and acceptation respectively, the
Figure 2: Schematic arrangement of involved atoms gain a full valence shell
an ionic lattice
A simple example of a compound which forms due to ionic
bonding illustrates the theory shown above. Sodium and chlorine combine to NaCl. Sodium
is in the first group of the periodic table and an element of period 3. For sodium it would be
energetically favourable to lose one electron to obtain a full n=2 valence shell. Chlorine,
however, is in group 7 and is an element of period 3. Therefore, for chlorine it would be
energetically favourable to gain one electron to obtain a full n=3 valence shell. Those two
desires are complementary; therefore the sodium transfers its high energy electron in the 3s
4 Two experiments in connection with intramolecular and intermolecular forces 2007 orbital to the chlorine atoms’ 3p orbital. As it is described later, this is in contrast to a
covalent bond in which both atoms involved in the bond have the same pull on the electrons
(i.e. the difference between the electronegativities of the atoms is equal). Therefore the
electrons are shared between both atoms. The electron transfer is visualised in a dot-and-cross
diagram in figure 3.
Figure 3: Schematic description of the electron transfer from sodium to chlorine when a ionic bond is formed
between the two atoms
The electrostatic attraction between the positively charged sodium ion and the negatively
charged chlorine atom draws them to each other, so that solid NaCl forms indeed a ionic
lattice as described above.
In a covalent bond electrons are shared between atoms (i.e. the probability of finding the
shared electrons is highest in a region directly between the nuclei of the atoms involved in the
bond). A purely covalent bond is virtually exclusively formed between two or more atoms of
the same element, because the electronegativity difference between the atoms has to be zero
in order to form a purely covalent bond.
A simple example of a molecule which
is formed by a covalent bond is a H2
molecule. Both atoms have the same
electronegativity, so that they attract
electrons with exactly the same amount.
However, both atoms would like to
obtain a full valence shell. Each
hydrogen atom has one electron in its
valence shell (1s) and would like to obtain one more electron to fill its valence shell. By
sharing their electrons, two hydrogen atoms can obtain this state. This is in contrast to the
formation of ionic bonds where the desires of the atoms are complementary. In the case of
ionic bonding both atoms involved need to gain electrons in order to reach the lowest energy
possible. Figure 4 visualises the concept of a covalent bond with the help of a dot-and–cross
diagram.
Figure 4: Schematic description of the formation of a
covalent bond between two hydrogen atoms
However, the most covalent bonds that form are covalent bonds which also have ionic
character. This happens when the electronegativity difference is approximately between 2 and
almost 0. This means that the probability of finding the shared electrons is higher nearer the
atom with the higher electronegativity. In a heteronuclear diatomic compound, for example,
the charge of the more electronegative atom of the molecule at an instance in time is to a very
high probability not zero but slightly negative, while the charge of the less electronegative
5 Two experiments in connection with intramolecular and intermolecular forces 2007 atom is slightly positive. Such a compound is said to be polar due the partially charges within
the electron. The atom with the higher electronegativity is said to have a partially negative
charge (δ-), while the atom with the lower electronegativity is said to have a partially positive
charge (δ+). However, the electrons are not transferred from one atom to the other. Because
the electrons are shared in the region between the atoms the distance between the nuclei of
the atoms involved in a covalent bonding is smaller than in ionic bonding. Therefore covalent
bonds are stronger than ionic bonds.
Whether a molecule has a permanent dipole or not is not only dependent on the involved
atoms, but also on the shape of the molecule. If a molecule is symmetrically arranged and
does involves only the same polar bonding pairs, the partially charges equalise each other,
and the molecule is not a polar molecule although it contains polar covalent bonds with ionic
character. An example of such a molecule is methane (CH4).
Bonding forces as ionic and covalent bonds are intramolecular forces. However, there are
also forces which act between molecules. Those forces are called intermolecular forces. In
comparison with ionic and covalent bonds those forces are very weak. However, they can
play a crucial role in chemistry and biochemistry. Physical properties for example as melting
point or boiling point are mainly a result of such intermolecular forces.
There are mainly two types of intermolecular forces: Van der Waals forces and hydrogen
bonds. In fact, there are actually more than two but they are of no concern in this experiment.
Van der Waals forces act between all molecules or ions, independently from their character.
Van der Waals forces come into existence also due to charge attraction between
molecules.They find their origin in the fact that electrons continuously, almost randomly
move within a molecule. Therefore the electron density is by chance higher at one end of the
molecule and lower at the other end of the molecule at random instance in time. Therefore an
instantaneous dipole is created at the instance in time. This instantaneous dipole is not to be
confused with the permanent dipole of a compound due to the electronegativity difference
between two atoms involved in a covalent bonding with ionic character.
However, this instantaneous dipole within one molecule is still an intermolecular force.
However, when such a molecule is considered to be within a region with other molecules, the
other molecules have instantaneous dipoles too. The instantaneous partially negative side of
one molecule is attracted by the instantaneous partially positively charged side of another
molecule. Those forces between the molecules are called induced dipole-dipole forces or
dispersion forces. Figure 5 visualises this concept.
Dispersion forces are not to be confused with permanent dipole-dipole forces. Permanent
dipole-dipole forces between molecules are forces which are the result of the attraction of
permanent dipoles of two or more molecules. Generally permanent dipole-dipole forces are
stronger than dispersion forces.
6 Two experiments in connection with intramolecular and intermolecular forces 2007 Figure 5: Dispersion forces between the molecules are due to induced dipoles of the molecules
Van der Waals forces are on their own very small, but when they act together they can have a
significant impact. Their strength increases with the size of the molecules between they act,
because the number of electrons, and therefore the potential induced dipoles increases.
Another intermolecular force is the so called hydrogen bond. Hydrogen bonds are strong in
comparison to van der Waals forces. However, in comparison with the covalent bonds, they
are very weak.
In hydrogen bonds, the electron of a hydrogen atom is virtually shared between two atoms of
different molecules. Thereby the hydrogen is covalently bonded to one of the atoms and the
hydrogen bond is formed between the hydrogen and the other atom. However, this does not
imply that the covalent bond between the first atom and the hydrogen is broken.
The two atoms involved in hydrogen bonds besides hydrogen itself are confined to being
oxygen, nitrogen and fluorine. This is due to the big electronegativity and to the atomic radii
of those elements. A hydrogen atom itself is a very distinct atom because it contains only one
electron. Therefore the atomic radius of a hydrogen atom itself is very small. The covalent
bond between one atom of the elements stated above and hydrogen leads into a strong
permanent dipole. Therefore atoms with very high electronegativity from other molecules are
attracted by the partially positive charge of the hydrogen atom.
However, only oxygen, nitrogen and fluorine are small enough to form a hydrogen bond with
the atom-hydrogen complex. Chlorine for example is too big to get close enough to the
hydrogen to form a hydrogen bond, although it also has a very high electronegativity.
Two experiments were carried out to test intermolecular forces. In the first experiment the
miscibility of substances were tested, while in the second experiment the dipole characters of
some molecules were under investigation.
The first experiment is based on thermodynamic changes within a system when two liquids
are mixed together. During the experiment two substances of the same volume are mixed
together. Beforehand, their temperatures are recorded and an average temperature of both
liquids together is calculated. After the mixing, the temperature of both liquids together is
measured. Both liquids are chosen in a way that it is guaranteed that they do not react with
each other. Due to intermolecular forces between the molecules of the two substances
7 Two experiments in connection with intramolecular and intermolecular forces 2007 however, the temperature after the mixing drops, increases or stays the same compared to the
average temperature of both liquids which is calculated before (more details in chapter 6).
The experiment is repeated with two more liquid pairs.
The second experiment tests the dipole character of several molecules in the liquid state.
Therefore a charged rod is held near a flow of the tested substance. Due to the electrostatic
charge of the rod a dipole is attracted towards the rod and is therefore deflected. The stronger
the dipole is, the stronger is the degree of deflection.
2.) Diagram
Figure 6: The set-up of the experiment to test the miscibility of several substances
8 Two experiments in connection with intramolecular and intermolecular forces 2007 Figure 7: The set-up of the experiment to test the dipole character of several molecules
3.) Method
3.1) The miscibility of liquids
3.1.1) Apparatus
Beaker, boiling tube, cotton wool, 2 measuring cylinder, 2 thermometers, ethanol,
cyclohexane, propanone, butane, distilled water
3.1.2) Saftey Notes
All the organic compounds used in the experiment were highly flammable – there should be
no naked flames in the laboratory. Most liquids also give off vapours which are irritating to
the skin, eyes and lungs. Eye protection and gloves should be worn all the time. Fume
cupboards should be used.
9 Two experiments in connection with intramolecular and intermolecular forces 2007 3.1.3) Procedure
The principle set-up of the experiment is shown in figure 6. Firstly, a boiling tube was placed
in a beaker. Afterwards, the empty space in the beaker was filled cotton wool to reduce heat
loss. 10 cm3 ethanol were measured with the help of a measuring cylinder. Subsequently, it
was given into the boiling tube in the prepared beaker.
Secondly, 10 cm3 cyclohexane were measured with the help of another measuring cylinder.
A thermometer was given into both, the ethanol and the cyclohexane to measure the
temperature. After the thermometers had been in the liquids for a few minutes the
temperatures were recorded.
Afterwards, the cyclohexane was given into boiling tube in the prepared beaker in which the
ethanol was given before and with the help of a thermometer the temperature of both liquids
together was recorded.
The experiment was repeated with two other pairs, namely propanone-butanone and
propanone-water.
3.2) Testing liquids for a permanent dipole
3.2.1) Apparatus
Burette, beaker, charged rod, ruler, distilled water, ethanol, propanone, hexane, cyclohexane,
ethyl ethanoate (ethyl acetate)
3.2.3) Saftey Notes
All the organic compounds used in the experiment were highly flammable – there should be
no naked flames in the laboratory. Most liquids also give off vapours which are irritating to
the skin, eyes and lungs. Eye protection and gloves should be worn all the time. Fume
cupboards should be used.
3.2.2) Procedure
The principle set-up of the experiment is shown in figure 7. Six burettes were each filled with
one of the following liquids: Distilled water, ethanol, propanone, hexane, cyclohexane and
ethyl ethanoate.
Firstly, the polarity of water was tested. Therefore the valve of the burette with water inside
was opened. Therefore water began to run out of the burette. The positively charged rod was
held next to the water, close to the opening of the burette. The water stream was deflected
towards the rod. With a ruler the degree of deflection was measured.
The same procedure was carried out with the other five substances in the same manner.
10 Two experiments in connection with intramolecular and intermolecular forces 2007 4.) Errors
4.1) The miscibility of liquids
¾ The smallest sub-division of the thermometer was 1º C. Therefore a systematic error
of +/- 0.5º C occurred for each temperature measurement.
¾ The smallest subdivision of the measuring cylinder which was used for the volume
measurements was 0.2 cm3. Therefore a systematic error of +/- 0.01 cm3 occurred
during each volume measurement.
¾ Parallax errors occurred during both temperature measurement and volume
determination
¾ Thermal energy maybe escaped through the boiling tube which was not covered with
cotton wool.
¾ Due to lack of time at the end of the experiment the final temperatures were read
almost immediately after placing the thermometer in the liquid. Therefore it could be
that the thermometer did not show the actual temperature of the liquid at that instance.
4.2) Testing liquids for permanent dipole
¾ The smallest sub-division of the ruler used for the deflection measurements was 1
mm. Therefore a systematic error of +/- 0.5 mm occurred during each deflection
measurement.
¾ A parallax error occurred during each deflection measurement, due to the angle on
which it was looked at the scale of the ruler
¾ There were five variables involved in the experiment, which influenced the degree of
deflection, whose values were arguably not exactly equal for all six measurements.
This is important because the experiment should not deliver an absolute degree of
deflection but a comparison of the degree of deflection between the tested substances.
The five variables were the following:
• Distance between the tip of the burette and measuring point (further away –
less deflection)
• Speed of flow (slower speed – greater deflection)
• Pressure of the substance in the burette (more pressure – higher speed)
• Distance between the tip of the rod and the flow (smaller distance – greater
deflection)
• Charge of the rod (higher charge – greater deflection)
5.) Results
5.1) The miscibility of liquids
Table 1 shows all the data that was measured during the experiment. As it is seen, the mixing
of cyclohexane and ethanol decreased the temperature of the system, while the mixing of
propanone and water increased the temperature in the system. The only consequence of
mixing propanone and butanone was a marginal decrease in the temperature of the system.
Therefore the temperature in the latter system stayed almost the same. As it seen in chapter 6
11 Two experiments in connection with intramolecular and intermolecular forces 2007 this marginal change in temperature is probably due to errors made during the experiment.
Table 2 shows the structural formulae of the substances used.
Liquid A Liquid B T (initial) A/Cº Ethanol Propanone Propanone Cyclohexane 22.1 Butanone 22.0 Water 22.0 T (initial) B/ºC 22.4 22.5 22.0 Average initial temperature/ Cº 22.25 22.25 22.00 Final temperature /Cº 20.00 22.00 27.5 ΔT/Cº ‐2.25 ‐0.25 +7.5 Table 1: Measured and calculated data obtained from the experiment: “The miscibility of liquids”
Substance Ethanol Formula CH3CH2OH Cyclohexane C6H14 Propanone CH3COCH2 Butanone C2H5COCH3 Water H2O Structural formula Table 2: Structural formulae of the substances used in the experiment: "The miscibility of liquids"
12 Two experiments in connection with intramolecular and intermolecular forces 2007 5.2) Testing liquids for permanent dipole
The degree of deflection for each tested substance and the structural formulae are shown in
table 3. Hexane and cyclohexane were not deflected at all, while ethyl ethanoate was
deflected the most (20 mm). Water was also deflected relatively strongly (11 mm). Ethanol
(3-4 mm) and propanone (6-7 mm) were only slightly deflected by the charged rod.
Substance Formula Distilled water H2O Ethanol CH3CH2OH Structural formula Deflection (mm) 11 3‐4 Propanone CH3COCH3 Hexane C6H14 6‐7 0 Cyclohexane C6H14 Ethyl ethanoate CH3COOC2H5 0 20 Table 3: Structural formulae and measured deflection of the substances used in the experiment: "Testing liquids
for permanent dipole"
13 Two experiments in connection with intramolecular and intermolecular forces 2007 6.) Conclusion/Evaluation
6.1) The miscibility of liquids
The mixing of water and propanone resulted in an increase of the temperature of the mixture
compared to the average temperatures of the separated liquid substances. This in turns means
that the reaction must be exothermic (i.e. releases thermal energy). Particles of both
substances must therefore arrange themselves in a way that is energetically favourable, i.e.
drops the energy of the two molecules. The structural formulae of both propanone and water
are shown in table 2. It is seen that there is a double bonded oxygen atom attached to carbon
2 in the ketone, which results in a strong polar bond between the carbon atom and the oxygen
atom. Due to the unsymmetrical arrangement of the molecule the molecule is a dipolar
molecule. Due to the higher electronegativity of the oxygen atom, the oxygen atom has a
partially negative charge, while the carbon 2 has a partially positive charge.
If the structural formula of water is considered, it is seen that it is also a very strongly dipolar
molecule. Thereby one oxygen atom is bonded to two hydrogen atoms by two polar bonds,
which result, due to the unsymmetrical arrangement of the molecule in the molecule being a
polar molecule. Due to the higher electronegativity of the oxygen atom in comparison to the
hydrogen atoms, it becomes partially negatively charged, while the two hydrogen atoms
become partially positively charged.
In view of the fact that both substances do not react chemically which each other in terms of
forming new chemical bonds or carrying out addition or substitution reactions respectively,
only a change in the strength of the intermolecular forces between the different molecules in
comparison to the intermolecular forces between the molecules of the same substance can be
the cause for the temperature change of the liquid. In view of the mixing being exothermic
the attraction between two molecules of different species must be greater than the attraction
between the molecules of the same substance.
In this case the reason for the greater attraction between the molecules of different species in
comparison to the attraction between molecules of the same substance is due strong hydrogen
bonds which from between the oxygen of the carbonyl group and a hydrogen atom of a water
molecule. To verify this statement both substances needed to be looked at on their own.
In liquid propanone the molecules are only attracted to each other by dispersion forces and
permanent dipole-dipole forces. However, due to the double bond, the dipole character
between the carbon atom and the oxygen atom of the carbonyl group is very strong. No
hydrogen bonds between propanone molecule form because there are no hydrogen-oxygen
(or nitrogen or fluorine) bonds in a propanone molecule.
In liquid water however, hydrogen bonds are formed between the water molecules. The
molecules arrange so that hydrogen bonds are formed between the oxygen atom of one water
molecule and one hydrogen atom of another one. Those hydrogen bonds are strong in
comparison with the dispersion forces which act between the propanone molecules.
14 Two experiments in connection with intramolecular and intermolecular forces 2007 However, due to the mixing being exothermic, the forces between a propanone and a water
molecule must be stronger than the hydrogen bonds between the water molecules. Due to the
strong dipole character of the bond of the carbonyl group, the partially positively charged
hydrogen atoms of a water molecule are more attracted by the partially negatively charged
oxygen atom of the carbonyl group than by the less partially negatively charged oxygen
atoms of another water molecule. The hydrogen bonds between the water molecules are
therefore broken and new hydrogen bonds between the hydrogen atoms of water molecules
and the oxygen atoms of the carbonyl groups of the propanone molecule are formed, which as
a result releases thermal energy which in turn increases the temperature of the system.
The mixing of propanone and butanone resulted in a small decrease in temperature. However,
propanone and butanone are both ketones. The only differences between both molecules are
an additional carbon and two hydrogen atoms attached to butanone in comparison with
propanone. The structural formulae of propanone and butanone are both shown in table 2.
Due to the bigger size of the butanone molecule, the dispersion forces between butanone
molecules are stronger than those between propanone molecules. However, they only differ
to a small extent. The measured temperature change was arguably bigger than the actual
temperature change. The temperature change should have been that small that it could not
have been measured in the experiment. The measured temperature change is arguably due to
errors which occurred in the experiment (see chapter 4).
The mixing of ethanol and cyclohexane resulted in a decrease of temperature. This means
that during to the mixing process of cyclohexane and ethanol energy was used.
In an ethanol molecule the oxygen atom and the hydrogen atom of the alcohol group are
bonded by a polar bond due to the electronegativity difference between the oxygen atom and
the hydrogen atom. Due to the unsymmetrical structure of an ethanol molecule, an ethanol
molecule is a polar molecule. When ethanol molecules come together hydrogen bonds are
formed between the alcohol groups of one ethanol molecule and a hydrogen atom the alcohol
group of another ethanol molecule. The intermolecular forces between the ethanol molecules
are therefore relatively strong.
Cyclohexane, however, is an alkane with six carbon atoms. Therefore no other bonds are
involved than carbon hydrogen bonds. As it is discussed in chapter 6.2 the electronegativity
difference between carbon and hydrogen is that small that it can be considered as a non polar
bond. Therefore cyclohexane is a non-polar molecule. The only intermolecular forces
between cyclohexane molecules are therefore dispersion forces.
As observed in the experiment the temperature drops when both substances, ethanol and
cyclohexane are mixed together. This is because the non-polar cyclohexane pushes the
ethanol molecules further apart, which in turns decreases the strength of the hydrogen bonds.
The action pushing the ethanol molecules apart requires energy, which is the reason why the
temperature decreases.
15 Two experiments in connection with intramolecular and intermolecular forces 2007 However, at the first glance this observation seems to contradict one fundamental principle of
chemistry, which is that interactions between molecules only occur when they are
energetically favourable. Why should the ethanol molecules move further apart when this is
energetically unfavourable? Or in other words, why do both liquids mix at all? Fats and
water, for example, do also not mix because the strong hydrogen bonds between the water
molecules prevent the mixture, so why should ethanol and cyclohexane?
The answer lies in a branch of physical chemistry called thermodynamics. Crudely described,
thermodynamics deals with the flow of energy. Thermodynamics is one of the most important
branches of chemical physics. It is the foundation of understanding the reasons of the
occurrence of any chemical reaction.
Thermodynamics usually regards the energy flow between a separated area und its
surroundings. In terms of thermodynamics this separated area is called a system, while the
surroundings are nothing less the rest of the universe. It follows:
“The Universe = the system + the surroundings”2
Thereby the system is not able to exchange matter with the surroundings and the other way
round. However, energy can be transferred between the system and the surroundings. The
system is usually a place in which chemicals are present or in which chemical reactions take
place.
If the system is in equilibrium with the surroundings the net movement of energy between the
system and the surroundings is zero. There is therefore a fixed amount of energy in the
system, which is mainly made up of bond energy between the molecules and heat, which is
nothing more than the kinetic and potential energy of the molecules within in the system.
“The energy contained in a system is called the internal energy (symbol U)”3.
If a chemical reaction takes place in the system, bonding energy gets free due to the reaction
in the majority of cases. The reaction is said to be exothermic. This energy firstly increases
the heat within the system. It is important to note that the reaction alone does not change the
internal energy of the system, only the form of energy changes from bonding energy to
thermal energy.
However, the temperature in the system is now higher than in the surroundings so thermal
energy flows out of the system into the surroundings to reach equilibrium again. The internal
energy of the system therefore decreases. The change in internal energy is therefore
expressed by difference of the final internal energy and the initial internal energy, the internal
energy the system had before the reaction took place. This relationship is given in
mathematically terms by the following equation:
For an exothermic reaction the change in internal energy is therefore negative.
16 Two experiments in connection with intramolecular and intermolecular forces 2007 However, the internal energy is not only influence by heat but also by the work done on
respectively by the system. Work can be defined in terms of pressure as the product of
pressure times the change in volume. If, for example, a chemical reaction in the systems
causes “a small increase in the volume (ΔV) of the”4 system “the walls of the” system “must
push against the constant pressure in the surroundings in order to expand”4. “This does work
on the”4 surroundings “and requires energy”4. “When a system expands, the work done on
the system equals –pΔV”5, because energy is needed to do work on the surroundings.
However, the work done by a system is dependent on two variables, the volume (V) and the
pressure (p). If the volume is constant, for example if the system is enclosed by an extremely
strong container, and heat is released due to a reaction the pressure inside the system
increases. However, if the volume is of the system is able to change, as for example in a
piston, the pressure does work on the surroundings.
The change in enthalpy (ΔH) is defined as in the following:
Where h is the heat added to the system.
If the pressure is constant (i.e. the change in pressure is zero), the change of enthalpy is equal
to the heat transferred. So:
At constant pressure.
If a reaction occurs the change in energy of the system is ΔU. “The first law of
thermodynamics states: Energy cannot be created or destroyed”6. So when energy is added to
the system from the surroundings, the energy of the surroundings must decrease by the same
amount. It follows that:
ΔU = -h
So the change of enthalpy at constant pressure is equal to ΔU.
The second law of thermodynamics states that “systems will change spontaneously from
states of lower probability to states of higher probability”7. An example illustrates the
statement of the law. If a box with 1000 coins in it, which all lie heads up, is shaken, there is
only a very low probability that all the 1000 coins still lie heads up, because there is only one
possibility this state can come into existence (all 1000 coins must lie heads up). However,
there is a much higher probability that the coins arrange so that the ratio of coins lying heads
up to coins lying tails up is 50:50, because there are much more possibilities how this state
can be achieved. The state of higher probability is therefore said to be more disordered.
“Entropy” (S) “is a measure of the disorder of a system”. The entropy of a system increases
when the matter or energy in the system spreads out or becomes more random in its
arrangement”8.
17 Two experiments in connection with intramolecular and intermolecular forces 2007 The change in entropy of a system is defined as:
Where h is the heat added to the system and T is the absolute temperature of the system. In
view of the first law of thermodynamics it can be stated that the sum of the entropies of the
system and of the surroundings must give the entropy of the universe, because the heat added
to system (h) is equal to the heat subtracted from the surroundings (-h), so that the energy of
the universe stays the same.
However, because the universe always changes to a state of higher probability or more
disorder and entropy is a measure of disorder the second law of thermodynamics can be
restated as the in the following:
“The total entropy of the universe always tends to increase, it never goes down”8.
So, if a reaction takes place in the system that requires energy (endothermic), it draws in
energy directly from the system and to maintain the equilibrium between the system and the
surroundings indirectly from the surroundings. In this case the entropy of the surroundings
decreases. Because of the second law of thermodynamics, which states that the entropy of the
universe always increases, the entropy within the system must increase to compensate the
entropy decrease in the surroundings. This increase is not due to an increase in heat, it is due
to an increase in the disorder of the arrangement of the molecules inside the system.
This is the explanation for the temperature decrease which occurred during the mixing of
ethanol and cyclohexane in the experiment. Thereby the system was the liquid, while the
surroundings were the rest of the universe. The endothermic reaction took energy from the
surroundings; this decreased the entropy of the surroundings. Therefore the entropy in the
system had to be increased in order to obey the second law of thermodynamics. This was
done by disordering the, due to the strong intermolecular forces, highly ordered structure of
the ethanol molecules with the help of cyclohexane molecules. As it is seen in chapter 6.2 the
dipole character of water is much stronger than that of ethanol. The intermolecular forces
between water molecules are therefore too strong to become disorder by fats, so that fats and
water do not mix. The entropy change of the universe would therefore at standard conditions
not be positive for the mixing of water and fats.
6.2) Testing liquids for a permanent dipole
Polar molecules contain covalent bonds between elements whose electronegativities differ.
Therefore the likelihood of finding an electron nearer the more electronegative atom is higher
than finding it nearer the atom of less electronegativity. The rod was rubbed with a fur in
order to make the rod carry an electrostatic charge. When the rod was placed near the liquid
stream, the rod attracted the partially charged parts of the atom of a substance, if there were
dipoles in the tested liquid. The stronger the dipole of the tested substance the higher was the
degree of deflection, supposed that the charge of the rod was the same.
18 Two experiments in connection with intramolecular and intermolecular forces 2007 Most compounds tested were organic substances. Organic substances are all composed of a
skeleton of carbohydrate chains. Accurately speaking, each carbon-hydrogen bond has a
small dipole (Electronegativity H: 2.1, Electronegativity C: 2.5). However, this dipole is that
small that it does not have any significant consequences to be considered in this experiment.
This dipole did not affect the deflection in a significant way.
However, most organic molecules contain, besides the hydrocarbon skeleton also functional
groups which are mostly composed of one or more atoms of oxygen and nitrogen (and
hydrogen) which are attached to a carbon atom of the carbon skeleton. Oxygen and nitrogen
have very high electronegativities, much higher indeed than those of carbon and hydrogen;
they are also able to form hydrogen bonds. In a bond with a carbon atom from the carbon
skeleton or a hydrogen atom those molecules highly polarise the bond. Therefore the
likelihood of finding the electrons is much higher near an oxygen atom or a nitrogen atom.
The partially charged sides of the dipole can therefore be attracted by the charged rod.
Ethyl ethanoate was deflected the most. This means the dipole of ethyl ethanoate is the
strongest of the substances tested. It is the strongest because two highly electronegative
oxygen atoms are bonded to one carbon atom. Therefore the carbon atom becomes strongly
partially charged. Therefore a strong dipole is formed, which was attracted by the charged
rod.
Water had the second strongest permanent dipole, followed by propane and ethanol. The
oxygen atom draws the two electrons from the hydrogen atom partially towards it due to its
high electronegativity. Therefore a strong dipole is formed. However it is weaker than the
dipole of ethyl ethanoate, because in ethyl ethanoate two oxygen atoms, one by a double
bond one by a single bond, are bonded to one carbon atom. Because the water molecule
forms two sigma bonds with two hydrogen atoms and because of the higher electronegativity
difference between hydrogen and oxygen in comparison with carbon and oxygen the water
molecule has a stronger dipole than propanone. The double bond between the carbon and the
oxygen of the carbonyl group of propanone consists of one σ-bond and one π-bond (a π-bond
is less strong than a σ-bond).
The carbonyl group of propanone is formed by polar bond. Because of the double bond
between the carbon and oxygen of the carbonyl group, the dipole of propanone is stronger
than the dipole of the oxygen atom and the hydrogen atom of the alcohol group of ethanol,
whose two atoms are only singly bonded. The higher electronegativity difference between
oxygen and hydrogen does not compensate the double bond of the carbonyl group in terms of
the strength of the dipole.
However, pure saturated carbohydrate chains (alkanes) like hexane or cyclohexane have no
functional group attached to them. Therefore there is no dipole in such a molecule, so that
they were not attracted by the positive charged rod.
19 Two experiments in connection with intramolecular and intermolecular forces 2007 7.) Evaluation
The experiment delivered the, from the structural formulae of the molecules deduced,
expected results. Therefore both experiments were successful.
8.) Bibliography
8.1) Notes
Note 1, 2, 3, 5, 6, 8: Clugston, Michael; Flemming, Rosalind. Advanced Chemistry. (2000).
Oxford University Press. pp. 616
Note 4, 7: Alberts, Bruce; Molecular biology of the cell; 2002; 4th edition; Garland Science;
pp. 1463.
8.2) Figures
Figure1:
http://www.xmission.com/~seldom74/chem1110int/ch03/electronegativity.jpg&imgrefur=
Figure 2:
http://www.bbc.co.uk/scotland/education/bitesize/standard/img/chemistry/propertiesofsubsta
nces/ionic/ionic_lattice.gif
20