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Canadian Chemistry Olympiad Lab Safety Guidelines
By its very nature, experimental chemistry is potentially dangerous. Although every
effort is made to identify and (where possible) eliminate risks, it is not possible to anticipate every hazard. Ultimately you, and you alone, bear the responsibility for your own personal safety in the laboratory. It is therefore imperative that you take the time to read and understand all the safety information and procedures outlined below. Much of lab safety, however, comes down to exercising practical common sense:
Exercise common sense at all times Always read the label on the bottle In case of chemical spills: wash immediately In case of fire: get out If in doubt: ask your lab supervisor Dispose of all chemical waste properly – not down the sink Always be aware of where you are and what you are doing
1. Dress Code:
Laboratory attire must be protective from head to toe, and easily removed in the event of
chemical spills. Appropriate clothing and safety equipment must be worn at all times while in the laboratory, even if you are finished your experiment: remember, others may still be working around you! Always wait until you are outside the lab to take off your lab coat and safety goggles.
Clothing:
Wear long pants and fully enclosed, durable footwear. Do not wear open-toed sandals or flip-flops, as they leave your feet exposed to any chemical spill. Similarly, shorts and skirts are not permitted.
Lab Coat: Wear a lab coat over your clothes, and keep it done up at all times in the laboratory.
Eye Protection: University policy requires that splash–resistant safety goggles must be worn at all times in the laboratory, even over prescription glasses, to protect against chemical splashes or flying objects. Safety goggles must be indirectly vented or non-vented, and comply with CSA Standard Z94.3-02.
Hair: Long hair should be tied back out of the way. Loose articles of clothing should also either be removed or secured so that they cannot fall into your experiment!
Hands: Gloves can provide protection against accidental chemical spills, but can also provide a false sense of security. They can even lead to accidental contamination of equipment and doors, or even cause experimental problems. Make sure that any gloves worn provide adequate protection for the specific chemical(s) you will be using. Remember that all disposable gloves are permeable to differing extents; nitrile gloves provide the greatest protection, but will not stand up to prolonged contact with concentrated acids and bases or halogenated solvents such as dichloromethane. • You must remove your gloves before:
* Handling any computers or instrumentation * Touching hair, skin, or clothing * Leaving the laboratory for any reason.
2. General Precautions:
• Smoking, eating and drinking are strictly prohibited in the laboratory. • Handle all chemicals with care, using rubber gloves where appropriate. . Always read the
label on a container to make sure it is both the right chemical and the right concentration. Thin, disposable gloves are not appropriate for use with halogenated solvents such as
chloroform; they have been known to leak on occasion. • When dispensing chemicals:
– always hold a container by the label to minimize accidental contact – always clean up chemical spills immediately; never leave a mess behind – always return chemicals promptly to the correct location – always make sure container lids are on properly
• Remove your gloves before using instruments and computers, or exiting the lab, in order to avoid possible chemical contamination.
• Make sure you have read, and understood, the correct procedures to follow for each operation such as weighing, making solutions, or filtering. Use of the correct procedure not only minimizes the risk of personal injury, but also ensures accurate, precise, and reliable data.
• Splashes of corrosive or toxic substances should be washed immediately from the skin or eyes with copious amounts of cold water. If a toxic or corrosive substance splashes your lab coat, take it off before the contamination can soak through to your clothing.
• Clean up any chemical spills immediately. Neutralise strong acids with bicarbonate, then wash with water. Soak up organic solvents using a cloth or towel, leaving it in a fume hood until the solvent evaporates.
• Report any broken thermometers to the lab technician immediately. • Always use the fume hood when instructed to do so. Never open the fume hood door
beyond the limits marked on the side frames. • Do not dispose of chemicals or broken glassware in garbage bins. Use the appropriate
container for all chemical waste, and always check the label on the container before emptying your waste into it.
• To avoid cuts, always handle glassware gently. If you break any glassware, ask for the broom and dustpan and sweep up the pieces as soon as possible. If you cut yourself on a piece of glassware, wash the cut immediately with cold water to avoid chemical contamination and apply pressure to stop further bleeding. Seek first aid from a lab demonstrator, technician, or the instructor in charge. Place broken glass in the correct bins:
– If the glass is contaminated with chemicals, place it in the green or black bins – If the glass is clean (washed at least three times), place it in the orange bins.
3. Fire Precautions: When using flammable solvents (e.g. ethanol, acetone, etc.) do not light matches or
Bunsen burners, and do not keep excess solvent in open beakers. Familiarise yourself with the location of all fire equipment, including the location of fire extinguishers, fire blankets, safety showers, fire alarms, and emergency exits.
In the event of a fire:
• If the fire is contained in a flask or beaker, cover the vessel to exclude oxygen. • If spilled organic solvent is burning on the floor or bench top, etc., sound the alert and
leave the lab without being asked to do so. • If someone near you is on fire, smother the flames with a lab coat or fire blanket. If the
safety shower is nearby, douse the victim. Do not wait for someone else to act.
Canadian Chemistry Olympiad Calculations, Errors, and Significant Figures
In any practical science, it is essential that all calculations and results should
take into account the errors and uncertainties associated with actual measurements. In this context, the error Ex in a measured value x is defined as the difference between the true and measured values: Ex = xtrue – x. While such measurement errors always exist, it is not always possible to determine their value. In fact, in most cases the best we can do is provide an estimate of the error, which we term the uncertainty; any value x therefore always has an associated uncertainty ±Δx, which must be accounted for in any subsequent calculations.
Determining the uncertainty, Δx:
• Derived or measured values always have an implied uncertainty if none is
stated, which is taken as ±1 in the last significant figure. Uncertainties always have the same units as the value for which they provide an error estimate. For example:
Mm = 242.13 ± 0.01 g/mol E = –163 ± 1 mV
A = 0.137 ± 0.001 AU C = 1.00 × 10–3 ± 1 × 10–5 mol/L
• When the arithmetic mean of a set of n repeated values is calculated, the uncertainty is based on the unbiased standard deviation, s. This is denoted on most calculators as either σ n–1 or sn–1. Note that a standard deviation, being an estimate of the error in the mean value, only has one significant figure; it is customary to report the first non-significant digit also, in order to avoid premature rounding in subsequent calculations. For example:
Example 1: m = 10.012, 10.005, 10.007, 10.004, 10.011 g
!
m = 10.0078 g, s = 0.00356 g, n = 5 ∴
!
m = 10.008 ± 0.004 g (n = 5) Note that, in this example, the mean value actually has fewer significant
figures than the individual readings as a direct consequence of the value of the standard deviation!
• Sometimes, an instrument is incapable of resolving the inherent variation between repeated measurements, resulting in a string of identical values with s = 0. In this case, the uncertainty is expressed as ±1 in the last figure of the instrument display. For example:
Example 2: E = 107, 107, 107 mV
!
E = 107 mV, s = 0 mV, n = 3
!
"E = 107 ± 1 mV (n = 3)
Error propagation with implicit uncertainties: When performing calculations using values that have an implied uncertainty,
it is essential to observe the correct rules for determining the uncertainty Δq in the final value, q, as this will determine the correct number of significant figures to which the result may be quoted.
• When adding or subtracting, the least significant digit in the result is
determined by the value with the fewest decimal places. • When multiplying or dividing, the least significant digit in the result is
determined by the value with the fewest significant figures. • When a calculation involves both types of operation (additive and
multiplicative), observe the order of mathematical precedence and determine the correct numbers of significant figures for intermediate values.
• It is always advisable to retain at least the first non-significant digit in
intermediate values in order to prevent premature rounding of the final result. Non-significant digits should be clearly denoted using either subscripting, underlining, or enclosing them in parentheses.
Example 1: calculate the molar mass of calcium carbonate
1 × 40.08 = 40.08 1 × 12.011 = 12.011 3 × 15.9994 = 47.9982 = 100.0892 ∴ Mm(CaCO3) = 100.09 g/mol
Notice how, in this example, the molar mass correctly has 5 significant figures, even though the atomic mass of calcium was only known to four significant figures! This is because the largest uncertainty in any of the masses is ±0.01 g/mol, so the error in the sum total lies in the second decimal place.
Example 2: calculate the mass of calcium carbonate required to make
100.0 mL of an 0.200 M solution:
!
C =m (g)
Mm
(g/mol)"V (L)#m = C "M
m"V
!
"m = 0.200mol
L#
100.0 mL
1000 mL/L#100.0892
g
mol
∴ m = 2.0018 g = 2.00 g of calcium carbonate
Notice how, in this example, the final result can only be given to 3
significant figures, even though the volume and molar mass are known to 4 and 5 significant figures, respectively. (Remember that leading zeroes before the decimal place are not counted towards the number of significant figures; we could clarify this by writing C = 2.00 × 10–1 M.)
• When calculating logs and antilogs, remember that the number of significant
figures associated with a value x determine the number of decimal places reported for log x. Similarly, the number of decimal places for a value y determines the number of significant figures reported for 10y.
Error propagation with calculated uncertainties: In analytical and physical chemistry, it is important to explicitly calculate
the effect of individual uncertainties on the overall uncertainty in the final result. In fact, the rules in the preceding section are based on the mathematical procedures one must use to do this. It is therefore important to evaluate the uncertainties in additive and multiplicative steps independently, as different rules apply.
• When adding or subtracting values with random errors, e.g. if q = ax + y
where a is a constant, the uncertainty Δq in q is:
!
"q = ± a"x( )2
+ "y( )2
Example: the uncertainty in the difference between two weights, each
measured to the nearest ±0.0001 g:
m = mfinal – minitial = 23.2476 – 21.1942 g = 2.0534 g
!
"m = 0.0001( )2
+ 0.0001( )2 = ± 0.00014 g
• When multiplying or dividing values with random errors, e.g. if q = ax × y
where a is a constant, the uncertainty Δq in q is:
!
"q = ±q"x
x
#
$ %
&
' ( 2
+"y
y
#
$ %
&
' (
2
Example: the uncertainty in the concentration obtained by dissolving
2.0534 g of CaCO3 in water and making to volume in a 100.0 mL volumetric flask
• The uncertainty in the mass from above is ± 0.00014 g • The molar mass from above is 100.09 ± 0.01 g/mol • The uncertainty in flask volume is ± 0.08 mL
C =
!
m
Mm"V
=
!
2.0543 g
100.0892 g/moL"
1000 mL/L
100.0 mL= 2.05157 × 10–1 mol/L
!
"C = ±C"m
m
#
$ %
&
' ( 2
+"V
V
#
$ %
&
' ( 2
+"M
m
Mm
#
$ %
&
' (
2
!
"#C = ±0.2051570.00014
2.0534
$
% &
'
( ) 2
+0.08
100.0
$
% &
'
( ) 2
+0.01
100.0892
$
% &
'
( )
2
∴ ΔC = ± 1.66 × 10–4 M
∴ C(CaCO3) = 0.2052 ± 0.00017 mol/L
Rounding values from calculations: Whether dealing with implied or actual uncertainties, it is essential to round
the final value correctly. Most textbooks describe the procedure adopted in computer science, and commonly implemented on calculators, where values are rounded up if the first non-significant digit is greater than or equal to 5, but rounded down if less than 5. Unfortunately, this leads to some bias in the results for cases where the non-significant portion of the number is exactly 5. In analytical chemistry, we therefore implement a third rule: if the non-significant portion of the number is exactly 5, the result is rounded to the nearest even number. This means that half the time the result will be rounded up, and half the time it will be rounded down.
Examples: q = 13.379512 – the non-significant digits are > 0.0005 so we round up q = 227.63439 – the non-significant digits are < 0.005 so we round down q = 1.1415 – the non-significant digits are = 0.0005; round to 1.142
q = 10.765 – the non-significant digits are = 0.005; round to 10.76
Page 1 of 4
Handling Volumetric Glassware
1. Obtaining Accurate Weights Modern analytical balances allow weights to be obtained to the nearest 0.1 mg
fairly easily. Proper technique is essential, however, if accuracy and precision are required. Most importantly, materials should be pure, dry, at room temperature, and both chemically and physically stable. Analytical balances are generally reliable, but can be prone to drifting or erratic readings if not used correctly. In general:
• Do not move the balance • Avoid touching weighing bottles with your bare hands • Avoid creating static charge on weighing bottles • Keep the balance pan free of chemicals • Always close the balance doors when weighing • Do not lean or tap on the bench while weighing • Always determine the mass transferred by difference
Procedure:
1. Calculate the mass required for your experiment. 2. Dispense approximately the mass required into a clean, dry weighing bottle using
a 2- or 3-decimal place top-loading balance. 3. Close the doors on an analytical balance, and set it to zero using the tare button. 4. Use crucible tongs or a folded slip of paper to place the weighing bottle on the
balance pan. * Do not handle weighing bottles with your bare hands or rubber gloves!
5. Close the balance doors, and record the reading once stable. 6. Use crucible tongs or a folded slip of paper to remove the weighing bottle, and
empty the contents into a suitable vessel. 7. Using crucible tongs or a folded slip of paper as before, re-weigh the weighing
bottle, remembering to close the balance doors before taking the reading. 8. Remove the weighing bottle and close the balance doors; leave the balance clean
and tidy for the next person.
Page 2 of 4
2. Quantitative Transfer The goal of quantitative transfer is to make sure that all of a solution is transferred
from one vessel to another. It is typically used when preparing standard solutions from solids, as it is often undesirable to transfer and dissolve solids directly into volumetric flasks. In general:
• Avoid spills and splashes • Keep glassware positioned where it won’t get knocked or fall over • Make sure you have everything you need before you start
Procedure:
1. Take note of the size of the volumetric flask; you do not want to use more than
about half this volume of solvent to dissolve the solid. 2. Dispense accurately the mass of solid required into a suitable size beaker. * You do not need to dispense exactly the right amount, as long as you know the
actual mass transferred as accurately as possible (i.e. by difference to 0.1 mg) 3. Add about a quarter to a half of the required volume of distilled water to the solid
in the beaker, and stir until the solid is fully dissolved. Funnel
Clamp
Volumetric flask
Decant downglass rod
Push last drop backinto beaker
4. Position a funnel in the neck of the volumetric flask using a ring clamp, leaving
an air gap between the stem of the funnel and the sides of the flask. 5. Decant the contents of the beaker down a glass rod into the funnel; use the glass
rod to push the final drop in the beaker’s spout back into the beaker. * Keep the glass rod in the beaker! 6. Use a wash bottle to rinse the walls of the beaker; decant the washings into the
flask as in step 5. Do this at least twice, making sure that you do not over-fill the flask (i.e. the solution should remain below the calibration mark on the neck.)
7. Rinse the glass rod into the funnel, then remove the flask from under the funnel. 8. Use a Pasteur pipette and rubber bulb to fill the flask to the calibration mark with
distilled water. Stopper the flask, and mix the contents by repeated inversion.
Page 3 of 4
3. Using a Transfer Pipette Transfer pipettes allow specific volumes of solution to be dispensed precisely and
accurately. They come in two basic types: ones marked ‘TD’ (‘to deliver’) and ones marked ‘TC’ (‘to contain’). The former are commonly used with aqueous solutions, and are calibrated to deliver the specified volume ± a small uncertainty when drained without forcing the final small volume of solution out. In general:
• Always rinse pipettes with a small volume of the solution to be dispensed • Never pipette by mouth – always use a pipette filler • Use a small beaker to hold the solution when filling the pipette • Keep the tip of the pipette in the solution when filling • Never shake a pipette to remove residual liquid after cleaning • Keep a clean paper towel or KimWipe™ handy to dry the pipette tip
Procedure:
1. Pour a small excess of the solution to be dispensed into a small clean, dry beaker. 2. To operate a three-valve pipette filler: use the top valve (1) to evacuate the filler;
place the filler on the end of the pipette without forcing it; use the middle valve (2) to draw solution into the pipette; use the lower valve (3) to drain it.
3. To rinse a pipette, draw a small volume of solution into the pipette (usually just
up to the lower part of the wide middle section.) Tip the pipette horizontally, slide the filler off the end, and let the solution run past the calibration mark while rotating the pipette. Drain the solution into a waste container. Rinse two or three times in total.
4. Fill the pipette by drawing the solution up past the calibration mark on the stem, but not so far that the solution enters the pipette filler.
5. Remove the pipette from the solution, and wipe any excess liquid from the outside of the pipette at the tip.
6. Holding the pipette over the beaker of solution, carefully drain the pipette until the bottom of the liquid meniscus just touches the calibration mark; if you drain to far, refill following steps 4 and 5, then try again.
7. Hold the pipette just inside the neck of the required volumetric flask, and drain the contents of the pipette into the flask under gravity. Lightly touch the tip of the pipette to the liquid surface in the flask to deliver the correct volume.
Page 4 of 4
4. Using a Burette Burettes allow controlled volumes of solution to be delivered precisely and
accurately. When operating the burette, make sure it is clamped in the lower half (about a third of the way up from the valve) in order to avoid stressing the glass. You should always interpolate the meniscus between scale markings to achieve the greatest precision. In general:
• Always rinse burettes with a small volume of the solution to be dispensed • Never reach overhead to fill a burette; lower it within the clamp, place the
clamp stand on a stool, or remove the burette from the clamp to fill it • Expel the air in the burette tip by opening the valve fully once filled; a full
height of liquid is needed to expel the air • Never shake a burette to remove residual liquid after cleaning • Keep a clean paper towel or KimWipe™ handy to dry the burette tip • Never waste time trying to get the initial reading to be exactly zero; this is
actually less accurate than recording a reading within the scale. Procedure:
1. Pour the solution to be dispensed into a clean, dry beaker. 2. Make sure the valve on the burette is closed (tap at right-angles to the burette),
and then pour a few millilitres of solution into the burette while rotating it. 3. Drain this solution into a waste beaker, and repeat step 2 one or two times more. 4. Fill the burette with solution up past the zero reading on the scale. Run a small
volume of the solution out in order to expel the air. 5. Reading a burette: take the reading with the liquid meniscus at eye level (move
the burette if necessary!). Interpolate the reading on a 50 mL burette between scale divisions to the nearest 0.01 mL:
“To deliver”
1st scaledivision
0.00 mL
0.10 mL
0.20 mL
0.30 mL
0.40 mL
0.50 mL
Reading is estimated as 0.18 mL
Canadian Chemistry Olympiad Calculations, Errors, and Significant Figures
In any practical science, it is essential that all calculations and results should
take into account the errors and uncertainties associated with actual measurements. In this context, the error Ex in a measured value x is defined as the difference between the true and measured values: Ex = xtrue – x. While such measurement errors always exist, it is not always possible to determine their value. In fact, in most cases the best we can do is provide an estimate of the error, which we term the uncertainty; any value x therefore always has an associated uncertainty ±Δx, which must be accounted for in any subsequent calculations.
Determining the uncertainty, Δx:
• Derived or measured values always have an implied uncertainty if none is
stated, which is taken as ±1 in the last significant figure. Uncertainties always have the same units as the value for which they provide an error estimate. For example:
Mm = 242.13 ± 0.01 g/mol E = –163 ± 1 mV
A = 0.137 ± 0.001 AU C = 1.00 × 10–3 ± 1 × 10–5 mol/L
• When the arithmetic mean of a set of n repeated values is calculated, the uncertainty is based on the unbiased standard deviation, s. This is denoted on most calculators as either σ n–1 or sn–1. Note that a standard deviation, being an estimate of the error in the mean value, only has one significant figure; it is customary to report the first non-significant digit also, in order to avoid premature rounding in subsequent calculations. For example:
Example 1: m = 10.012, 10.005, 10.007, 10.004, 10.011 g
!
m = 10.0078 g, s = 0.00356 g, n = 5 ∴
!
m = 10.008 ± 0.004 g (n = 5) Note that, in this example, the mean value actually has fewer significant
figures than the individual readings as a direct consequence of the value of the standard deviation!
• Sometimes, an instrument is incapable of resolving the inherent variation between repeated measurements, resulting in a string of identical values with s = 0. In this case, the uncertainty is expressed as ±1 in the last figure of the instrument display. For example:
Example 2: E = 107, 107, 107 mV
!
E = 107 mV, s = 0 mV, n = 3
!
"E = 107 ± 1 mV (n = 3)
Error propagation with implicit uncertainties: When performing calculations using values that have an implied uncertainty,
it is essential to observe the correct rules for determining the uncertainty Δq in the final value, q, as this will determine the correct number of significant figures to which the result may be quoted.
• When adding or subtracting, the least significant digit in the result is
determined by the value with the fewest decimal places. • When multiplying or dividing, the least significant digit in the result is
determined by the value with the fewest significant figures. • When a calculation involves both types of operation (additive and
multiplicative), observe the order of mathematical precedence and determine the correct numbers of significant figures for intermediate values.
• It is always advisable to retain at least the first non-significant digit in
intermediate values in order to prevent premature rounding of the final result. Non-significant digits should be clearly denoted using either subscripting, underlining, or enclosing them in parentheses.
Example 1: calculate the molar mass of calcium carbonate
1 × 40.08 = 40.08 1 × 12.011 = 12.011 3 × 15.9994 = 47.9982 = 100.0892 ∴ Mm(CaCO3) = 100.09 g/mol
Notice how, in this example, the molar mass correctly has 5 significant figures, even though the atomic mass of calcium was only known to four significant figures! This is because the largest uncertainty in any of the masses is ±0.01 g/mol, so the error in the sum total lies in the second decimal place.
Example 2: calculate the mass of calcium carbonate required to make
100.0 mL of an 0.200 M solution:
!
C =m (g)
Mm
(g/mol)"V (L)#m = C "M
m"V
!
"m = 0.200mol
L#
100.0 mL
1000 mL/L#100.0892
g
mol
∴ m = 2.0018 g = 2.00 g of calcium carbonate
Notice how, in this example, the final result can only be given to 3
significant figures, even though the volume and molar mass are known to 4 and 5 significant figures, respectively. (Remember that leading zeroes before the decimal place are not counted towards the number of significant figures; we could clarify this by writing C = 2.00 × 10–1 M.)
• When calculating logs and antilogs, remember that the number of significant
figures associated with a value x determine the number of decimal places reported for log x. Similarly, the number of decimal places for a value y determines the number of significant figures reported for 10y.
Error propagation with calculated uncertainties: In analytical and physical chemistry, it is important to explicitly calculate
the effect of individual uncertainties on the overall uncertainty in the final result. In fact, the rules in the preceding section are based on the mathematical procedures one must use to do this. It is therefore important to evaluate the uncertainties in additive and multiplicative steps independently, as different rules apply.
• When adding or subtracting values with random errors, e.g. if q = ax + y
where a is a constant, the uncertainty Δq in q is:
!
"q = ± a"x( )2
+ "y( )2
Example: the uncertainty in the difference between two weights, each
measured to the nearest ±0.0001 g:
m = mfinal – minitial = 23.2476 – 21.1942 g = 2.0534 g
!
"m = 0.0001( )2
+ 0.0001( )2 = ± 0.00014 g
• When multiplying or dividing values with random errors, e.g. if q = ax × y
where a is a constant, the uncertainty Δq in q is:
!
"q = ±q"x
x
#
$ %
&
' ( 2
+"y
y
#
$ %
&
' (
2
Example: the uncertainty in the concentration obtained by dissolving
2.0534 g of CaCO3 in water and making to volume in a 100.0 mL volumetric flask
• The uncertainty in the mass from above is ± 0.00014 g • The molar mass from above is 100.09 ± 0.01 g/mol • The uncertainty in flask volume is ± 0.08 mL
C =
!
m
Mm"V
=
!
2.0543 g
100.0892 g/moL"
1000 mL/L
100.0 mL= 2.05157 × 10–1 mol/L
!
"C = ±C"m
m
#
$ %
&
' ( 2
+"V
V
#
$ %
&
' ( 2
+"M
m
Mm
#
$ %
&
' (
2
!
"#C = ±0.2051570.00014
2.0534
$
% &
'
( ) 2
+0.08
100.0
$
% &
'
( ) 2
+0.01
100.0892
$
% &
'
( )
2
∴ ΔC = ± 1.66 × 10–4 M
∴ C(CaCO3) = 0.2052 ± 0.00017 mol/L
Rounding values from calculations: Whether dealing with implied or actual uncertainties, it is essential to round
the final value correctly. Most textbooks describe the procedure adopted in computer science, and commonly implemented on calculators, where values are rounded up if the first non-significant digit is greater than or equal to 5, but rounded down if less than 5. Unfortunately, this leads to some bias in the results for cases where the non-significant portion of the number is exactly 5. In analytical chemistry, we therefore implement a third rule: if the non-significant portion of the number is exactly 5, the result is rounded to the nearest even number. This means that half the time the result will be rounded up, and half the time it will be rounded down.
Examples: q = 13.379512 – the non-significant digits are > 0.0005 so we round up q = 227.63439 – the non-significant digits are < 0.005 so we round down q = 1.1415 – the non-significant digits are = 0.0005; round to 1.142
q = 10.765 – the non-significant digits are = 0.005; round to 10.76
Experiment 6: Multistep Synthesis 1. Part B: Acetylation of p-Aminophenol 1
Experiment 6
Multistep Synthesis 1. Part B: Acetylation of p-Aminophenol
Pre-lab Assignment
1. Review nucleophilic acyl substitution reactions. 2. Write the mechanism for acetylation of 4-aminophenol. 3. Calculate how many millimoles of each reagent will be used in today’s reaction. 4. Read about thin-layer chromatography (pp 37-40).
Introduction
Acetylation of Arylamines
In the synthesis of phenacetin, the introduction of the N-acetyl group is an essential step in conferring the physiological activity (analgesic and antipyretic) for which the final product is important. More often however, in other syntheses, primary and secondary amines are converted to their acetyl derivatives as a 'protective' measure to reduce or moderate their susceptibility to oxidative degradation and electrophilic substitution or to 'block' their reaction with another functional group or reagent. Even in the present exercise, this role is played to a certain extent by acetylation at this stage in the synthesis. Both arylamines and phenols undergo acetylation in the presence of acetic anhydride or acetyl chloride. It is evident therefore, that, since p-aminophenol contains both of these functional groups, the acetylating agent and other experimental conditions will have to be carefully chosen and controlled if selective N-acetylation (to the exclusion of O-acetylation) is to be achieved. Resolution of this problem (which arises quite frequently in organic syntheses) depends upon their being at least a slight difference in reactivity between the two types of functional groups. Because oxygen is more electronegative than nitrogen, amines are in general, stronger bases and nucleophiles than are phenols and therefore tend to react faster in the nucleophilic substitution reactions characteristic of acid derivatives.
For the acetylation of arylamines, acetic anhydride is usually favoured over acetyl chloride for several reasons. Acetyl chloride is less desirable because it is a disagreeable reagent to manipulate and half of the amine escapes acetylation through conversion to its hydrochloride salt (by the hydrogen chloride liberated in the acetylation). Moreover, for this specific synthesis, the reactivity of acid chlorides is so great that even the weakly nucleophilic phenolic group would also undergo acetylation to a significant extent.
Acetylation of arylamines proceeds smoothly and completely with acetic anhydride, either alone or in acetic acid solution. Indeed the lower reactivity of acetic anhydride (relative to the acid chloride) is such that its rate of hydrolysis is sufficiently slow to permit acetylation of amines to be carried out in aqueous solutions. Phenolic substances however, unlike amines, cannot be acetylated satisfactorily in an aqueous medium since there is no significant difference in the reactivity of the acetic anhydride toward the OH group of the phenol over that of water. Indeed, you will recall from the preparation of aspirin that the rate of reaction of phenols with acetic anhydride is sufficiently slow to require the addition of a few drops of concentrated sulfuric acid as a catalyst.
Experiment 6: Multistep Synthesis 1. Part B: Acetylation of p-Aminophenol 2
Therefore, p-aminophenol is converted to its N-acetyl derivative, p-hydroxyacetanilide, by warming an aqueous suspension with acetic anhydride until solution takes place. On cooling, the product crystallizes from solution in almost quantitative yield. Experimental
Suspend 1.0 g of 4-aminophenol in 4 mL of distilled water in a 10-mL round-bottomed flask. Add 1.2 mL of acetic anhydride. Stir the mixture in a 55°C water bath until the solution becomes clear.
Continue heating for 10 minutes then cool the reaction mixture in an ice bath until crystallization of the product appears to be complete. Collect the crystals by suction filtration in a Hirsch funnel. Wash the product with ice-cold water and dry under suction. Check the purity of your product by TLC using ethyl acetate - glacial acetic acid (9:1, v/v) solvent system. Run IR and 1H NMR spectra. Report
1. Report the yield and melting point of your product. 2. Attach the spectra, assign peaks and report spectral data. 3. Calculate and report the Rf values for 4-aminophenol and 4-hydroxyacetanilide. 4. Amides undergo hydrolysis to carboxylic acids under both acidic and basic conditions. Draw the
important resonance contributors for the resonance-stabilized cation formed during acid-catalyzed hydrolysis of 4-hydroxyacetanilide.
HN
C
CH3
O
HO
H2O
N H2
HO
+ + CH3C OOH
C H3
C
O
CC H3
O
O
warm
Experiment 6: Multistep Synthesis 1. Part B: Acetylation of p-Aminophenol 3
1H NMR Spectrum of 4-Hydroxyacetanilide (300 MHz, CDCl3)
13C NMR Spectrum of 4-Hydroxyacetanilide (75.7 MHz, CDCl3)
170 160 150 140 130 120 110 100 90 80 70 60 50 40 30 20
25.03
76.65
77.08
77.51
117.51
117.94
130.71153.12167.94
TLC
THIN-LAYER CHROMATOGRAPHY (TLC)
Chromatographic techniques are used extensively in the organic laboratory for both qualitative separations and quantitative analysis. A thorough understanding of the various types of chromatography is essential. Thin-layer chromatography (TLC) is used to determine the purity of a compound, to evaluate how far a reaction has proceeded, and to analyze the composition of a mixture, while column chromatography is used to physically separate components of a mixture. High-pressure liquid chromatography (HPLC) can be used for both qualitative and quantitative analysis.
All chromatographic techniques are based on a similar principle: components of a mixture can often be differentiated by exposure to two competing phases. In TLC, the stationary phase is a polar adsorbent, such as silica gel or alumina, which has been coated onto a glass or plastic plate. The mobile phase is an organic solvent or mixture of solvents. The liquid moves up the plate by capillary action. Column chromatography is similar to TLC, with a stationary phase of silica gel or alumina and a mobile liquid phase, but differs in that the liquid travels down the column. A polymer-coated adsorbent is the stationary phase in HPLC. The mobile liquid phase in HPLC is forced through a small diameter column by a pump at high pressure. Thin-Layer Chromatography (TLC) is the separation of moderately volatile or nonvolatile substances based upon differential adsorption on an inert solid (the stationary phase) immersed in an organic solvent or solvent mixture (the mobile phase). The components are distributed between the stationary phase (usually silica gel or alumina) and the solvent depending upon the polarities of the compound and solvent. The compounds are carried up the plate (ascending chromatography) at a rate dependent upon the nature of the compounds and the solvent.
Figure 1. Adsorption equilibrium for molecules between adsorbent and solvent.
Compounds are separated by adsorption chromatography based upon differential attachment of molecules to the adsorbent and the polarity of the solvent used for the separation as shown in Figure 1. Polar compounds are strongly attracted to and held by a polar adsorbent. Nonpolar compounds are held weakly. When a nonpolar solvent is passed through the adsorbent, nonpolar compounds are released easily, but polar compounds are retained. When a moderately polar solvent is passed through the adsorbent, both nonpolar and polar compounds are released, but the nonpolar compounds move faster because there is still an attraction between the polar compounds and the polar adsorbent. In general, nonpolar compounds move faster than polar compounds for TLC on silica gel or alumina.
TLC
Polarity of Solvents Used for Adsorption Chromatography Solvents are rated according to their polarities. Common solvents are listed below from most polar to least polar.
Polarity Solvent Formula Dielecetric constant, ε
Most polar
Least polar
methanol ethanol acetone methylene chloride ethyl acetate diethyl ether toluene cyclohexane hexanes
CH3OH CH3CH2OH CH3COCH3 CH2Cl2 CH3COOCH2CH3 CH3CH2OCH2CH3 C6H5CH3 C6H12 C6H14 (isomers)
32.6 24.6 20.7 9.08 6.02 4.34 2.38 1.965 1.89
If a polar compound moves too slowly in a nonpolar solvent, switching to a more polar solvent will cause
the compound to move faster. If a nonpolar compound moves too fast on TLC, switching to a less polar solvent will cause the compound to move slower. The polarity of the solvent system can also be varied by mixing miscible solvents to give the desired separation.
Calculating Rf Values
The distance traveled by each component is expressed as a rate or retardation factor (Rf). Rf values are calculated by dividing the distance traveled by a component by the distance between the origin and the solvent front (distance traveled by the solvent). Hence all Rf values will fall between 0 and 1. Rf values are measured from the origin here the initial spot was applied to the centre of the spot.
In general, Rf values on silica gel or alumina are decreased by decreasing the polarity of the solvent and increased by increasing the solvent polarity. The main objective in TLC analysis is to obtain separation of components of a mixture. A solvent is selected to give a range of Rf values. If a single solvent fails to give adequate separation, a solvent mixture should be used.
TLC
Resolution In TLC, it is highly desirable to avoid overlapping of spots in the developed chromatogram, just as it was desirable to avoid overlapping peaks in GC. In GC, resolution can be improved by lowering the column temperature, decreasing the carrier gas flow rate, and using a different column or longer column. In TLC, resolution can be increased by varying the nature of the eluting solvent or by changing the adsorbent. An example of TLC that shows the effect of solvent polarity is given below. For the separation of benzaldoxime (B) and p-tolualdoxime (T), hexane failed to move the compounds significantly, but ethyl acetate moved both compounds too fast, giving little separation. In this case, a mixed solvent system containing both hexane and ethyl acetate gave better results, moving both compounds about midway up the plate, while also achieving different Rf values for B and T. The exact percentage of a mixed solvent system is best determined by trial and error. In this case, the optimal mixture was 70% hexane/30% ethyl acetate.
Adsorbent
The most frequently used adsorbents are silica gel (Si02) and alumina (A1203), which are coated onto a plastic or glass support. TLC plates may be purchased from commercial vendors or they may be homemade. Commercial plates have adsorbents bound to a plastic, glass, or aluminum sheet as a backbone. The binder is calcium sulfate. The adsorbent is very uniform, about 0.1-0.2 mm thick. Homemade plates are usually made by dipping glass microscope slides into mixtures of adsorbent and binder in aqueous solution. Homemade plates are not usually as uniform in thickness as commercial plates. However, since TLC is generally used qualitatively rather than quantitatively, it is not so critical that the thickness of the layer be that uniform. Coated microscope slides are wiped clean on one side and the coated side is left to dry. These plates may be used directly for TLC. Commercial plates may be sized appropriately using scissors or a glass cutter.
Sample Application
Dots (in pencil) are marked uniformly along one of the narrow edges of the plate near the bottom at intervals of about 1-2 cm. Solutions of samples are spotted individually at each of the marks and the solvent is allowed to dry. The diameter of the residue from a spotted solution should be kept as small as possible to minimize diffusion effects. Capillary tubes or Pasteur pipettes that have been drawn out in a flame are used as applicators. The very small openings in these drawn-out tubes permit application of very small amounts of solution.
TLC
Development The development chamber can be a beaker covered by a watch glass or a jar with a screw-top cap. The
chamber should be equilibrated before use by placing the developing solvent and a piece of filter paper in the chamber and capping the chamber. The chamber should stand undisturbed for 5-10 minutes to saturate the chamber with solvent vapours. The plate is put into the developing chamber after making certain that the level of the solvent is below that of the applied compounds on the plate. As the chromatogram is developed, the solvent advances up the plate through capillary action. Unless coloured compounds are being separated (such as dyes in inks), the components cannot be visualized. The plate is left in the developing chamber until the level of the solvent is about 1-2 cm from the top of the plate. Then the plate is removed from the chamber, the solvent line is marked immediately with pencil, and the plate is allowed to dry.
Visualization
Because most organic compounds are not coloured, the spots must be visualized using a UV light, an iodine chamber, or an indicator spray. Certain substances, such as aromatic compounds, absorb ultraviolet light and appear as purple spots. Commercial plates often contain a fluorescent dye that gives a light green background when exposed to UV radiation. Compounds that do not absorb UV light must be viewed in another way, such as placing the plate in an enclosed chamber containing solid iodine crystals. Most organic substances form coloured complexes with iodine. After a few minutes, organic compounds on the developed plate begin to appear as brown spots. The effect is reversible over time after removal from the iodine chamber, so spots should be circled. Another option for visualizing spots is to spray the plate with a reagent that will cause development of a colour. A commonly used spray is anisaldehyde-HzS04, which gives coloured spots for alcohols and certain other compounds. Procedure The overall TLC procedure is summarized in Figure below.
(a) A solution is spotted at the origin on a TLC plate, using a drawn capillary tube.
(b) The plate is placed in a development chamber containing solvent. The level of solvent is below the origin. (c) The solvent front is marked immediately upon removing the developed TLC plate from the development
chamber. The dried TLC is visualized and spots are circled.
1
An Approach For Tackling Organic Structural Determination Problems
Introduction A regular challenge to organic chemists is that of structure elucidation – determining the atom connectivity and stereochemistry of a molecule by interpretation of spectroscopic data. Typically one will have some molecular information from several sources, such as infra-red, nuclear magnetic resonance and mass spectrometry experiments. The first two techniques are those most routinely used in order to deduce the structure of an organic molecule. When presented with spectroscopic information about a particular molecule, it is essential to ADOPT A LOGICAL STEPWISE APPROACH in order to deduce its structure. The aim of this document is to outline a useful method that can be applied to ANY problem of this nature. The Approach Typically the following information is available for a compound of unknown structure: 1) Molecular formula (or elemental percentages from combustion analysis) 2) Infra-Red (IR) data (either as a spectrum or as important absorbances) 3) Nuclear Magnetic Resonance (NMR) data – usually proton (1H) and carbon (13C)
(either as spectra or as important absorbances) FIRSTLY – study the molecular formula of the compound and calculate its DEGREE OF UNSATURATION (DOU). This gives you important information about the molecules composition before analyzing ANY spectroscopic data. Example: DOU in a molecule of formula C8H9NO: i) Write formula of alkane with same number of carbon atoms = C8H18 ii) Ignore O in the formula iii) For each N, remove one hydrogen from C8H9NO – leaves C8H8 iv) Calculate DOU by subtracting number of H atoms in formula from number of H
atoms in alkane with same number of carbon atoms, then dividing by 2: (18-8)/2 = 5 DOU.
2
If the number of DOU ≥ 4 then it is very likely (but not definite) that a benzene ring is present in the molecule. If there is one DOU present then there may be one ring (e.g. a cycloalkane) or one double bond (e.g. an alkene) present. SECONDLY – study the INFRA-RED DATA supplied. Focus particularly on the following regions if presented with a spectrum: i) 1650 – 1850 cm-1 – the C=O (carbonyl group) strongly absorbs within this range.
It is not essential to try and identify which particular carbonyl group is giving rise to the absorption, this will generally become apparent from any NMR information.
ii) 3300 – 3600 cm-1 – an alcohol O-H bond appears within this range (a very
strong, broad absorbance). An N-H bond appears in the same area, but generally exhibits a weaker and sharper absorbance.
iii) 2100 – 2500 cm-1 – the characteristic triple bond region (due to C≡C or C≡N). THIRDLY – study the NMR DATA provided – starting with the 1H (proton): i) Count the number of distinct peaks in the spectrum (disregarding spin-spin
splitting) – this will clarify how many different “types” of hydrogen are present. ii) Look at the integration – the numbers next to each peak are indicative of how
many hydrogens are giving rise to the peak. iii) Consider the chemical shift of each peak – providing information about which
“types” of hydrogen are present, e.g. aromatic hydrogens generally appear in the range δ 6.5 – 8.0, whereas alkene hydrogens tend to resonate between δ 4.5 – 6.5.
iv) Analyze the spin-spin-splitting – is there any? (if not, then there will be no
hydrogens on adjacent carbon atoms in the molecule). On the other hand, a triplet-quartet splitting pattern may well indicate the presence of an ethyl group, -CH2CH3.
If supplied with 13C (carbon) NMR information then the number of different signals will make the number of “different” carbon atoms apparent. Also, the chemical shifts are again characteristic of different “types” of carbon atom. IT IS OFTEN ONLY NECESSARY TO LOOK AT THE CARBON NMR DATA IN TERMS OF STRUCTURAL CONFIRMATION – it is possible to glean much more detail about the molecule from the 1H NMR spectrum.
3
An Example Compound X: C10H12O2 Infra-Red Spectrum:
13C NMR Data: δ 20.8, 35.1, 64.8, 126.5, 128.5, 128.8, 137.8, 170.8.
4
1H NMR Spectrum:
What is the structure of X??
Two excellent websites for practice structure elucidation problems can be found at: http://www.chem.ucla.edu/webspectra http://www.nd.edu/~smithgrp/structure/workbook.html Prepared By: Dr. Andy Dicks Department Of Chemistry, University Of Toronto Tel: (416) 946-8003