Experiment 2. Separation of Liquids by Distillation.
References:
Eaton, Laboratory Investigations in Organic Chemistry
INTRODUCTION:
Experiment 2 illustrates a technique used for the separation of organic liquids – distillation. It
will also give you experience in infrared spectroscopy.
This experiment is a distillation of two different liquids – an alcohol (methanol, ethanol or 1propanol) and another liquid (acetone, diethyl ether or water). It will be up to you to discover
which two liquids are present in the mixture.
BACKGROUND:
Distillation Techniques
Distillation is a common method used to separate and purify liquids. The process consists of
heating a liquid to its boiling point, condensing the vapours, and collecting the condensate.
Figure 1 shows several stages in the evolution of distillation equipment in the organic laboratory.
Up until recent years, equipment based on the retort design was common. However, when
performing microscale experiments, a long distillation path must be avoided. With small sample
quantities, there is too much glassware (nooks and crevices, etc.) in a large scale apparatus and
too many opportunities to lose the entire sample. Microscale equipment based on the retort
design (A) is expensive and awkward, so we use the Hickman still design (B).
Figure 1: The Evolution of Distillation Equipment
General Principles
If a liquid is kept in a sealed contai
container,
ner, some molecules escape from the liquid's surface into the
space above it. When equilibrium is established, the number of molecules escaping the liquid
equals the number of molecules in the vapour that strike the liquid surface and stick. The
molecules in the vapour also strike the walls of the container and exert a pressure, defined as the
vapour pressure of the liquid. If the temperature of the liquid is raised, a greater number of
molecules escape to the vapour phase until equilibrium is once again established; the vapour
pressure of the liquid increases with increasing temperature. This is demonstrated for 4 common
liquids.
Figure 2:: Vapour Pressure
Pressure-Temperature
Temperature Diagram of 4 liquids
The boiling point of a liquid is that temperature at which the vapour pressure of the liquid is
equal to the pressure of its surroundings. If the flask that contains the liquid is open to the
atmosphere, the boiling point will be the temperature at which vapour pressure of the liquid is
equal to atmospheric pressure. The vapour pressure of a pure liquid rises steadily as the
temperature is increased, until the boiling point is reached (Figure 2).
Simple Distillation
1. A Mixture of Ideal Liquids
If a mixture of two misciblee liquids with different boiling points is heated to boiling, the vapour
will not have the same composition as the liquid; it will be richer in the more volatile component.
component
Consider
Figure 3 (top of the following page), which depicts the behaviour of mixt
mixtures
ures of Carbon
tetrachloride (C) and Toluene (T), two miscible volatile liquids with boiling points BPC and BPT,
respectively. The lower of the two solid curves represents the boiling points of different mixtures
of C and T. The upper curve represents the composition of the vapour which is in equilibrium
with the liquid at its boiling point. At 100% C or 100% T the curves meet, because boiling pure
C (at BPC) can have only pure A vapour in equilibrium with it; the same applies to pure T (at
BPT).
For example,
e, a mixture of 60% T and 40% C will boil at 93 oC. Moving horizontally at 93 oC
over to the upper vapour curve, we find that it intersects at a vapour composition of 39% T and
60% C (point A).
As the distillation proceeded, C would be selectively removed from the liquid. The composition
of the liquid would change gradually to 100% T. The boiling point of the liquid would gradually
rise to BPT; at the same time, the composition of the distillate (vapour) would gradually change
to 100% T. Thus, in a simple distillation
istillation of a mixture, the first material to distill (sometimes
called the first cuts or first fractions) will be rich in the most volatile or lower
lower-boiling
boiling
components and the last material (last cuts or fractions) will be rich in the less volatile or higherh
boiling components.
Figure 3: Liquid--Vapour Composition Diagram of Toluene-Carbon
Carbon Tetrachloride
This behaviour is true for "ideal liquids" with no intermolecular interactions, and is
approximated by many organic mixtures.
2. Liquids that Form Azeotropes
Some liquid mixtures, instead of forming ideal solutions with liquid
liquid-vapour
vapour composition
diagrams typified by Figure 3, form either minimum or maximum boiling mixtures. Typical
composition diagrams for such mixtures are shown in Figure 4. Any mixt
mixture
ure with a composition
exactly equal to that of CMIN or CMAX will distill at a single constant temperature, just as if it
were a pure liquid. Ethanol and water form a constant boiling mixture (95.6% ethanol, 4.4%
water) with a minimum boiling point of 78.2o C (lower than that of pure ethanol, 78.5o C, or pure
water, 100o C). Formic acid and water also form a constant boiling mixture (22.6% formic acid,
77.4% water) with a maximum boiling point of 107.2 o C (higher than that of formic acid,
100.8oC or water). Such constant boiling mixtures are called azeotropes (Greek: to boil
unchanged).
Liquid-Vapour
Vapour Diagrams of 2 Different Azeotropic Mixtures;
Figure 4:: Liquid
A) Ethanol and Water, B) Formic Acid and Water.
Mixtures with a composition to the lleft of CMIN, in Figure 4A (or CMAX in Figure 4B), can be
separated into pure ethanol (or formic acid in 4B) and the constant boiling mixture, but pure
water cannot be obtained from such a mixture by distillation. Conversely, mixtures with a
composition to the right of CMIN, in Figure 4A (or CMAX in 4B), can be separated into pure water
and the constant boiling mixture, but can never furnish pure ethanol/formic acid by distillation.
Fractional Distillation
It was noted from Figure 3 that even the first drop of distillate (fraction collected) obtained was
not pure, but rather a mixture, containing mainly the lower boiling point compound (carbon
tetrachloride). If these first fractions were to be combined and redistilled, the first vapour to be
condensed would be even richer in carbon tetrachloride. Repetition of this process (vaporization,
condensation, and revaporization) could eventually lead to isolation of pure carbon tetrachloride.
Similar redistillation of the higher
higher-boiling
boiling fractions could lead to isolation of pure toluene in the
final fractions. Clearly this repeated redistillation is a laborious process.
Figure 5:: Close
Close-up
up of the Fractionating Section of a Hickman Still
The fractionating column is a device for increasing the efficiency of this re
redistillation
distillation process. It
consists of a vertical column packed with some inert material, such as glass beads or glass
helices, or provided with some other device (indentations are used in our apparatus – see Figure
5) for increasing the surface upon which the vapour may condense. As the hot vapours rise
through the column, they condense and flow back down the column. The condensate, as it hits
the lower, hotter portions of the column, is revaporized, and the m
more
ore volatile components
proceed up the column once again. This process is repeated numerous times in the fractionating
section of the Hickman still, and the distillate will consist of the lowest
lowest-boiling
boiling component in a
nearly pure form. Figure 6 illustrate
illustrates the process graphically.
Figure 6: Liquid-Vapour Diagram, illustrating the Principles of Fractional
Distillation
An A-B mixture (Figure 6) with composition C1 boils at temperature TC1 and the vapours enter
the column at that temperature. If they condense in the column, the condensate will have the
composition C2. This vaporizes near the bottom of the column at temperature TC2, producing
vapours with composition C3. These may condense further up the column at TC3; vaporization
now gives vapour with composition C4, etc. If the column contains enough surface area for many
successive vaporizations and condensations (an ideal distillation column), the first distillate that
comes over will be nearly pure 100% A with a boiling point close to that of pure A. This will
continue until all of A is removed, after which B will start distilling, and the temperature will rise
rather abruptly to the boiling point of B. In practice, distilling columns are not 100% efficient,
but columns have been designed which can separate liquids that boil as little as 2oC apart.
Thermometer Correction
Generally speaking, thermometers may be of two types: total immersion or partial immersion. If
the thermometer is not immersed totally or to the immersion line on the thermometer, the
temperature reading will be slightly lower than the true boiling point of your distillate. The stem
correction may be calculated according to the formula:
Tc = To + F x L(To - Tm)
where Tc = corrected temp.,
To = observed temp.,
Tm = mean temp. of exposed stem,
L = length of exposed mercury column in degrees,
F = correction factor (for mercury in pyrex, a value for F of 0.00016 is used).
Although you are not expected to perform this correction, keep in mind that your observed
temperature will be lower than the true boiling point, especially for higher temperatures.
The Melting Point Apparatus
Several types of apparatus used to determine melting points are shown below. Types A to C are
older versions and use liquids as the heat transfer agents. In this laboratory you will use
apparatus D which comes in two versions: one for use with a capillary tube and the other for use
with a microscope cover slide. You should use the version with capillary tu
tubes. Make sure that
o
you use a high temperature (250 or 360 C) thermometer!
PRE-LAB
LAB PREPARATION
PREPARATION:
Read the experimental procedure so that you are prepared for the lab and you understand the
safety and disposal information for the chemicals you are using in this experiment.
1. How does vapor pressure determine the boiling point of a mixture (answer in the form of a
definition for boiling point)? What is the point of doing a fractional distillation?
2.. (a) Draw the structures of acetone, diethyl ether and water.
(b) Using the IR instrument, how would you differentiate these three liquids (include values)?
3.. Starting with a mixture of 50% carbon tetrachloride and 50% toluene (see Figures
Figure 3 AND 6),
and using an IDEAL fractional distillation column, what would be the % composition of toluene
in the first distillate (first fraction collected)?
EXPERIMENTAL PROCEDURE
PROCEDURE:
THIS LAB WILL BE PERFORMED IN PARTNERS
PARTNERS,, but each person must run their own
IR.
Safety and Disposal
osal Data for Compounds to be Distilled and Test Reagents.
Reagents
Compound
Acetone
Mol. Wt. (g/mol)
58.08
Safety and Disposal Data
Irritant. Highly Flammable. Dispose in Organic Waste.
Ceric Ammonium
Nitrate
548.23
Irritant. Contact with combustible may cause fire.
Wear gloves when handling.
Diethyl Ether
74.12
Irritant. Extremely flammable. May form explosive
peroxides. Dispose in Organic Waste.
3,5-Dinitrobenzoyl
230.56
Corrosive. Causes burns. Avoid all contact. Wear
chloride
gloves when handling.
Ethanol
46.06
Highly Flammable. Dispose in Organic Waste.
Methanol
32.04
Highly Flammable. Toxic by inhalation, in contact
with skin and if swallowed. Dispose in Organic Waste.
1-Propanol
60.09
Highly flammable. Irritant. Dispose in Organic Waste.
Equipment Used
All the equipment needed to perform this experiment should be in one of the three drawers in the
top row of your workstation, one of the drawers underneath your fumehood or the common
counter for your TAs group.
However the microkit (microscale glassware) should be obtained directly from your TA. At the
end of the experiment, return your CLEAN complete microkit and get the TA to initial the return
of your microkit.
1. Setting up the Apparatus
a) Obtain approx. 8 mL of ‘Michele’s Mix’ and transfer it to the 10 mL cylindrical vial.
b) Equip the vial with a couple of boiling chips and the Hickman
distillation head (see diagram to the left).
c) Start heating your apparatus. NOTE: Do not let the temperature
rise too rapidly. If you distill too quickly, you will not get a good
separation of your distillates.
d) Once boiling commences, the rate of heating should be lowered
to the point where the mixture is still boiling, but the temperature
increases no more than 2-3 °C/min. This is achieved by turning
the dial on the hot plate back until the light turns off, and then
slowly increase from there.
e) When the condensate starts collecting in the still, adjust the heat
so that liquid is slowly removed from the vial (it should take ~10
min. to fill the still). Once the boiling point of your mixture goes
above 60 °C, you may need to use the Aluminum Accessories when heating to push the vapour
up into the distillation head. Tip: It is very difficult to tell when there is a liquid in the
distillation head, try tilting the apparatus to see if any liquid moves.
2. Collecting the Fractions
a) When the well in the Hickman still is between ½ full and full (see Figure 5), remove the
fraction using a bent plastic pipette (Tip: this can be tricky! Make sure you put a significant bend
in the pipette!) and transfer it to a small test tube. You may need to remove the thermometer to
do this, is so, replace it as soon as you collect the fraction. If your fraction has a low boiling
point (< 65 °C), cover the test tube with parafilm to prevent evaporation.
b) Record the temperature range of the fraction in a table (see table below). The temperature
range is the temperature at which the fraction started to distill (ie collect in the distillation head)
and the temperature at which you collected the fraction using the pipette. Leave enough room in
your table for up to 10 fractions. Generally the ending temperature of one fraction will be the
starting temperature of the next fraction, except when all of the higher boiling liquid is gone. At
that time there will be an increase in the temperature before the second liquid starts to collect.
Tip: The best way to do this in partners: one person do the collecting of the fractions and
parafilming the tubes if necessary, while the other person is in charge of recording the
temperature ranges.
1st
fraction
2nd fraction
3rd fraction
4th fraction
5th fraction
6th fraction
Etc.
Temp. Range
Alcohol Test
(Colour)
c) Graph the fractions number (x-axis) against the middle temperature (from the range) for that
fraction.
3. Testing the Fractions for Alcohol
a) Add 1-2 drops of your distillate fraction to 0.5 mL of the yellow Ceric Ammonium Nitrate
solution (DO NOT add the yellow Ceric Ammonium Nitrate solution TO your fraction!).
b) If an alcohol (ROH) is present, a red colour should develop due to the complex
(NH4)2Ce[(OR)(NO3)5]. Record the colours and its brightness in the table. i.e. light orange or
dark red
4. Characterizing your Non-Alcohol
Based on your alcohol tests (step 3) and the temperature range of your fractions, you should be
able to distinguish the purest fractions of your 2 liquids.
a) Take the purest non-alcohol fraction and obtain an IR of your distillate (label pertinent peaks
– those above 1600 cm-1).
b) Based on your IR and boiling point of your distillate (Diethyl Ether 35°C, Acetone 57°C, and
water 100°C), do you have acetone, diethyl ether or water?
Remember: the observed BP is lower than the actual BP, especially at higher temperatures.
5. Characterizing the Alcohol
a) Take your purest alcohol fraction and put 0.3 mL of your ‘unknown’ alcohol in a 5 mL
conical vial.
b) Add 0.1g of 3,5-dinitrobenzoyl chloride1 and attach a drying tube packed with calcium
chloride.
c) Heat the vial on a hot plate for 5 minutes, swirling periodically to wash the powder from the
walls of the vial, then let the vial cool to room temperature.
d) Add distilled water (not regular tap water) to the 3 mL mark and shake vigorously until you
get a precipitate. You want to break up any large chunks of solid (Be careful not to mash your
powder into the bottom corner of the vial). Cool in an ice bath.
1
This is stored in a desiccator. It must be returned to the desiccator after use to prevent loss of activity due to air
moisture.
e) Separate the solid from the solution by vacuum filtration on a Hirsh funnel.
The vacuum filtration is done as follows (refer to the diagram
on the left). A Hirsch funnel is fitted to a filter flask with a
neoprene adapter. A disk of filter paper exactly the correct size
to cover all the holes in the funnel is placed in the funnel and
moistened with some of the solvent (in this case, water). The
filter flask is then joined to the vacuum inlet (using thick
orange or black tubing) and a vacuum is applied. When the
filter paper is drawn down tightly to the funnel, filtration of the
solution is begun. During filtration, it is important to break the
vacuum at either end of the hose connection before turning off
the vacuum. This will prevent loss of powder in the funnel due
to the pressure “suck back”.
f) Wash the precipitate twice with 5% sodium bicarbonate (while in the Hirsh funnel, add the
sodium bicarbonate and keep filtering) and allow the crystals to dry as much as possible.
g) Fill a melting point tube with your solid by thrusting the open end into the solid several times.
In order to work the plug of solid material down to the sealed end of the tube, vigorously tap the
sealed end on the table or lightly draw a file across the tube held loosely in the hand. Repeat the
procedure until the tube contains a 3 mm column of densely packed powder is in the bottom.
h) Place the capillary in the melting point apparatus, turn on the power and allow the hot-stage
temperature to rise fairly rapidly to within 15-20oC below the expected melting point of the
compound. However, during the determination of the actual melting point range, the temperature
should not rise more rapidly than 2 or 3o per minute. A finite time is required to transfer heat
from the hot-stage both through the walls of the capillary tube and throughout the mass of the
sample. If the hot-stage is heated too quickly, its temperature will rise several degrees during the
time lag required for the melting process to occur; this results in an observed range higher than
the true one.
i) Determine the melting point (MP) of the solid derivative (recorded as a RANGE, from when
the solid starts to melt until it’s done, I know, the name “melting POINT” is deceiving!) and this
will tell you which alcohol you have. The derivative MPs are as follows: Methanol 108°C,
Ethanol 93°C, 1-Propanol 73°C. Based on your derivative test and the boiling point of the
distillate (Methanol 65°C, Ethanol 78°C, 1-Propanol 97°C), which alcohol do you have?
Note: Since this is a problem based lab, you will be required to include a conclusion of your
findings.