Donald Dalton
ID#: 12320627
1/26/15
Chem. 244, 072
Lab Report 1: Purification of Solids
I. Introduction
This experiment was designed in order to demonstrate the very widely used
purification/extraction techniques of recrystallization and sublimation. By exploiting the physical
properties of certain organic chemicals such as solubility and enthalpy of sublimation,
purification of impure solids may be performed. The first two purifications were simple
dissolution reactions of benzoic acid and acetanilide in an excess of water. The third was a
dissolution and centrifugation of naphthalene in ethanol. The fourth purification was the
sublimation/deposition of 1,4-dichlorobenzene. All starting reagents were considered impure.
For a given dissolution reaction, there is a quantity that reports the number of moles of
solute that will completely dissolve in a given solvent at STP. This is called the standard molar
solubility (moles/L) and it describes the number of moles of solute that will saturate a liter of a
particular solvent. Once the molar solubility has been exceeded, solute will precipitate out of
solution at a rate equal to dissolution of additional solute in accordance with Le Chatelier’s
principle. For a solvation reaction the equilibrium constant K sp describes the relative
concentrations of the dissolved species once equilibrium has been achieved [1]. A high K sp
indicates high solubility while a low K sp indicates low solubility. Given that this is an equilibrium
value, the only external factor that can affect it is temperature, which is why there are different
Ksp values for different reaction temperatures. In general, for a dissolution reaction, the K sp and
molar solubility of a given compound will usually always increase with increasing temperature.
Likewise as the temperature of a dissolution reaction is decreased, the K sp and molar solubility
of a given compound will also almost always decrease.
Parts A/B/C of this experiment operate under the same general theory, an impure solid
is solvated in a quantity of solvent in which it will not completely dissolve. Heat is introduced to
increase the equilibrium quantity of solute that is dissolved until a miscible solution has been
formed. The solution is then allowed to cool at a slow rate. As the temperature drops so does
the molar solubility of the dissolved species and crystals begin to precipitate out of solution. As
this is a spontaneous process, thermodynamics dictates that the formation of the most stable
products will occur. Pure crystal lattices will produce the lowest value (most negative) of Gibbs
free energy change and impurities will not be incorporated into the growing crystals [2]. This is
an example of a reaction in which the enthalpy change dictates the free energy change and
reaction spontaneity despite an unfavorable change in entropy, Equation 1.
Part D utilizes a different technique for the purification of an organic solid, specifically
sublimation. Sublimation is the process by which a given chemical undergoes a phase change
from a solid directly to a gas forgoing the liquid phase. Recall there are three states of matter:
solid, liquid, and gas. Transitions between these three states are achieved by the input or
withdrawal of energy from a system. As a chemical in a given phase is heated, it will absorb
energy that is observable in the form of a temperature change. Once a certain threshold has
been reached, a phase transition will begin in which all energy is utilized to change states and
a temperature change will not be observed. The amount of energy required to provoke a phase
change is dependent on the number of moles of the chemical and the particular phase
transition. A constant called the enthalpy of formation, or heat of formation (kJ/mole) describes
the amount of energy per mole that is required to transition from one phase to another [3]. The
two most commonly used constants are the heat of fusion (melting) and the heat of
vaporization (boiling). However, if it were of interest in this experiment, heat of sublimation
would be utilized for energy calculations, as this is the process occurring. This type of
purification can only be implemented on chemicals that will readily sublime with the intention of
leaving impurities behind in the solid or liquid phase. Introducing heat into the system allows
for the sublimation of the desired pure chemical out of the impure solid. The gas then
undergoes deposition back into the purified solid state as it contacts a cold finger condenser.
General reactions are reported in Equations 2/3/4/5 for parts A/B/C/D respectively.
Equation : Gibbs free energy [2].
Equation 2: Solvation of impure benzoic acid in water.
Equation 3: Solvation of impure acetanilide water.
Equation 4: Solvation of impure naphthalene in ethanol.
Equation 5: Sublimation of impure 1,4-dichlorobenzene.
II.
Chemical Table
The following chemicals were utilized in experiment parts A/B/C/D.
Table : Chemicals utilized for parts A/B/C/D [4].
III.
Mechanism
The mechanisms for the parts A/B/C are best understood by analysis of the reaction
kinetics. In order to demonstrate the solubility theory at work, generalized reaction rate and
equilibrium expressions are shown below in Equations 6/7/8. Solvents will be assumed
pure and in excess and all reagents/products are assumed first order.
Equations 6/7: Dissolution equilibrium constant and reaction rate constant.
Equation 8: Arrhenius reaction rate constant [5].
As stated previously, equilibrium values for dissolution reactions typically
always increase with increasing temperature. Upon inspection of the reaction rate
equation, the temperature dependence can also be seen through the Arrhenius
rate constant K. Due to the temperature value in the denominator of the negative
exponential expression, larger values of temperature increase the value of K and
therefore increase the reaction rate.
The mechanisms for part D are best described in terms of the thermodynamic
events occurring during sublimation. As stated previously, sublimation is the
process by which a solid is converted directly into a gas forgoing the liquid phase.
The amount of energy required to perform sublimation can be calculated with the
number of moles and the enthalpy of sublimation constant seen below in
Equation 9.
Equation 9: Energy of sublimation equation [3].
Utilizing the energy from the heating mantle, the sample of 1,4dichlorobenzene readily sublimes into the gas phase. The gas then contacts the
cold finger condenser where it undergoes deposition free of impurities.
IV.
Results
The results from parts A/B/C/D are presented in Table 2. Percent extraction and percent
error were calculated via Equations 10 and 11. Percent error is reported as a range based
upon the observed high and low melting points relative to theoretical values in Table 1.
Table : Percent extraction, theoretical vs. observed melting points, and percent error.
Equations 10/11: Percent extraction and percent error calculations [4].
V.
Discussion/Conclusion
Upon inspection of the results in Table 2, there are several notable pieces of data that
merit discussion. Firstly, the extraction percentages range from [11–47]%. Without knowing
the percent composition of the impure chemicals it is not possible to make accurate
assumptions on the success of each recovery. However, if all of the starting reagents
contained relatively equal amounts of impurities, a trend of increased percent extraction is
observed implying increasing precision. The highest percent extraction occurred in the
sublimation reaction at 47%. This is significantly larger than any of the other percentages
reported for the first three recrystallizations. This provides strong evidence for the
effectiveness of sublimation purification when this technique is a viable option.
The second array of experimental data regards the melting points and is indicative of
the purity of the extracted chemical. Utilizing theoretical melting points for each chemical, a
percent error range was calculated based upon the high and low values of observed
melting points. Specifically, the temperature at which the first bit of solid turns to liquid and
the temperature at which the last bit of solid turns to liquid. The percent error ranges from a
low of [0.245 - 0.4]% for benzoic acid to a high of [3.7 – 4.9]% for naphthalene. Another
important feature of the melting point data is the ranges of melting points taken. Pure
chemicals tend to have sharp narrow ranges while impure chemicals tend to have broader
ranges. Table 2 indicates that all of the ranges calculated were within < 1.5 °C.
An important aspect of experimental analysis is reflection on potential sources of error
that may affect the data. Several mistakes were made throughout the execution of certain
procedures that may have affected the percent recovery as well as the melting point purity.
When doing a hot filtration through filter paper it is important that the solution remain at a
high temperature and that rapid cooling is avoided in order to ensure pure crystal growth.
During parts A/B, since they were performed simultaneously, the solutions were rapidly
poured into the filter funnel. As a result, the solution temperature dropped resulting in a
decrease in solute solubility. Crystals therefore began to form in the filter paper rather than
passing through to the filtrate. Crystal formation in the filter paper leads to a decreased
amount of captured pure solute and results in a decreased percent recovery. In the future
this may be avoided by keeping the solution on a hot plate and filtering smaller volumes of
solution at a time.
Another source of error was not weighing the vials in parts A/B in which the collected
samples were left to dry. This is crucial because the weight will not be accurate until wet
samples are given several days to dry. In order to compensate for this error, the samples
were removed from the vials once dry and weighed out. However, some of the sample
stuck to the interior of the vial and could not be removed. Again, this is a potential source of
error resulting in decreased recovery.
Avoiding mistakes such as those described will provide more accurate data in future
experiments.
VI.
Questions
1) Describe the steps for the recrystallization of biphenyl from ethanol.
a. Take biphenyl sample in centrifuge tube, add ethanol, and place in heat bath.
b. Allow the sample to heat until it forms a miscible solution.
c. Remove and slow cool in air for 10 min., then cool in ice bath for 10 min.
d. Centrifuge the sample for 1-2 min.
e. Pipet liquid out of tube leaving crystal, add cold ethanol, centrifuge again.
f. Pipet liquid out again, remove from tube, let sample dry.
2) Two steps can be left out in the recrystallization of biphenyl, which ones, when?
a. The two steps that could be left out of the recrystallization of biphenyl are the two
centrifugation steps. If the sample is allowed to cool at a slow controlled rate,
there will be adequate crystal growth free of impurities. Biphenyl has a density of
roughly 1 g/mL while ethanol has a density of 0.79 g/mL. So while the
centrifugation helps the separate the less dense ethanol from the more dense
naphthalene, this is not absolutely necessary, as this process will naturally occur
via gravity. Centrifuging is not necessary when dealing with compounds of
grossly different densities, as they will separate spontaneously.
3) How do you pick a good solvent for recrystallization? List four criteria [6].
a. The solute should be near insoluble in the solvent at room temperature and
extremely soluble in the solvent at high temperatures.
b. Impurities should either be very soluble in the solvent at low temperatures or
near insoluble in the solvent at high temperatures. Either of these scenarios
would allow filtration and separation of the pure chemical from the impurities.
c. The solvent should be relatively inert to the solute. No reactions between the two
should occur at any temperature.
d. The solvent should be easily removable by a number of different methods from
the extracted crystals i.e. washing/evaporation.
4) You have a grossly impure compound that refuses to recrystallize and does not sublime.
How can you initially purify the compound enough so that you can finish the purification
process?
a. This question is tough to answer because there is no information given on the
physical properties of the impurities that are contaminating the sample. However,
there are a number of different methods that could be implemented to try and
remove some of the bulk impurities before proceeding to the recrystallization
process. If it is known that there are acid/base differences between the pure
chemical and impurities, one might try and perform an acid/base extraction in
order to exploit solubility differences in protic and aprotic solvents. If there are
significant differences in polarity between the pure chemical and the impurities,
one might try and perform thin-layer or flash chromatography. If there are
significant differences in boiling points, one may try a simple distillation or a
fractional distillation. If there are significant differences in densities, one might try
and completely dissolve the sample and then separate by centrifugation. Again, it
all depends on the physical properties of the pure chemical and the impurities.
Once you have this information there are many different methods that can exploit
one of these properties and facilitate a purification process.
VII.
References
[1] "Solubility equilibrium". Retrieved January , 2015 Available:
http://en.wikipedia.org/wiki/Solubility_equilibrium
[2] "Gibbs free energy". Retrieved January, 2015 Available:
http://en.wikipedia.org/wiki/Gibbs_free_energy
[3] "Enthalpy of Sublimation". Retrieved January , 2015 Available:
http://en.wikipedia.org/wiki/Enthalpy_of_sublimation
[4] Organic Chemistry Lab Manual. Peter A. Wade Ed. 4, Vol. 1, 2014.
[5] "Arrhenius equation". Retrieved January, 2015 Available:
http://en.wikipedia.org/wiki/Arrhenius_equation
[6] "Recrystallization Technique". Retrieved January, 2015 Available:
https://www.erowid.org/archive/rhodium/chemistry/equipment/recrystallization.html