UNIT 13A- 359489
RISK ASSESSMENT FOR THE THREE PRACTICALS
Hazards:
Corrosive chemicals (sodium hydroxide solution, hydrochloric acid,
ethanoic acid)
Potential for chemical spills/splashes
Broken glassware
Risk Controls:
Wear personal protective equipment (PPE) - safety goggles, lab coat,
closed-toe shoes
Work in a fume hood when handling acids/bases
Add acid to water, never water to acid
Have spill kits and shower/eyewash stations available
Use equipment like pipette bulbs, never mouth pipette
Dispose of chemical waste properly
No food/drink in the lab
Secure long hair/loose clothing
Receive proper training on chemical handling
Have a chemical hygiene plan
Know location of safety data sheets (SDS)
Potential Health Hazards:
Skin/eye burns from corrosive chemicals
Respiratory irritation from inhaling vapors
Cuts from broken glassware
Chemicals:
1. Sodium Hydroxide Solution (0.1 M)
Corrosive
Causes severe skin burns and eye damage
Inhalation of mist/vapors can cause respiratory irritation
Wear protective goggles, gloves, clothing
Have eyewash and safety shower readily available
, Use in a fume hood
2. Hydrochloric Acid (0.1 M)
Corrosive, toxic
Causes severe skin burns and eye damage
Inhalation of vapors can cause coughing, choking
Wear protective goggles, gloves, clothing
Use in a well-ventilated area or fume hood
Have neutralizing materials available
3. Ethanoic Acid (0.1 M)
Corrosive, flammable
Causes skin and eye irritation
Inhalation of vapors may cause dizziness, drowsiness
Wear goggles, gloves
Use in a fume hood
Keep away from heat/sparks/flames
4. Ammonia Solution (0.1 M)
Corrosive, toxic
Causes skin/eye burns
Vapors highly irritating to eyes/respiratory system
Wear goggles, gloves, protective clothing
Use only in a fume hood
METHODS FOR THE PRACTICALS
PRACTICAL 1: Determination of the Ionization Constant of a Weak Acid (Ethanoic
Acid)
Prepare a standard solution of sodium hydroxide (NaOH) of known concentration.
Prepare a solution of ethanoic acid (CH3COOH) of unknown concentration.
Pipette a known volume of the ethanoic acid solution into a conical flask.
Add a few drops of phenolphthalein indicator to the flask.
Fill a burette with the standard NaOH solution.
Titrate the ethanoic acid solution with the NaOH solution, swirling the flask
continuously, until the solution turns pale pink.
,Record the initial and final burette readings to determine the volume of NaOH
used (titre value).
Repeat the titration process multiple times (e.g., 4-5 runs) to obtain concordant
values of the titre.
Calculate the mean titre value from the concordant readings.
At the half-equivalence point, measure the pH of the solution using a pH meter or
indicator paper.
Repeat the pH measurement for the half-neutralized solution in subsequent runs.
Calculate the average pH value of the half-neutralized solution.
Use the average pH value to determine the ionization constant (Ka) of ethanoic
acid using the following equations: pKa = pH; Ka = 10^(-pKa)
PRACTICAL 2: Study of Buffer Solution Properties
Prepare a buffer solution by mixing a weak acid (e.g., acetic acid) and its
conjugate base (e.g., sodium acetate) in appropriate proportions.
Measure the initial pH of the buffer solution using a pH meter or indicator paper.
Divide the buffer solution into several portions (e.g., 10 cm^3 each).
To one portion, add a few drops of a strong acid (e.g., hydrochloric acid, HCl) and
record the final pH.
To another portion, add a measured volume (e.g., 1 cm^3) of the strong acid and
record the final pH.
Repeat the process by adding increasing volumes of the strong acid (e.g., 2
cm^3, 4 cm^3, 5 cm^3, 10 cm^3) to separate portions of the buffer solution,
recording the final pH each time.
Repeat steps 4-6 using a strong base (e.g., sodium hydroxide, NaOH) instead of
the strong acid.
Observe and record the changes in pH for each addition of the strong acid or
base.
Analyze the results to study the buffering capacity of the buffer solution and its
ability to resist pH changes upon the addition of small amounts of strong acids or
bases.
PRACTICAL 3: Determining pH curves for acid-alkalis titrations
RESULTS FOR PRACTICALS (results for practical 3 are in the Excel file
PRACTICAL 1
, Trial 1st 2nd 3rd 4th
run run run run run
Initial reading/
0.45 9.60 0.65 7.70 0.15
cm3
Final reading /
25.50 35.35 26.25 33.85 25.95
cm3
Titre (volume
25.05 25.75 25.60 26.15 25.80
used)
Mean titre (V) /
38.5
cm3
V/2 / cm3 19.25
25.775/2= 12.8875= 12.90 to 2 d.p
pH of half-neutralised ethanoic acid 12.2
solution (1st run) 1
pH of half-neutralised ethanoic acid 12.1
solution (2nd run) 7
pH of half-neutralised ethanoic acid 12.1
solution (3rd run) 6
12.1
Average pH value
8
pKa= pH Ka= [H+]
pKa= 12.18 [H+] = 10-pH
[H+]= 6.61 x 10-13
Ka= 6.61 x 10-13
PRACTICAL 2
Results
10cm3 of buffer with HCl:
Few drops 1 cm3 2 cm3 4 cm3 5 cm3 10 cm3
(3)
Start pH 5.07 5.07 5.08 5.08 5.08 5.08
,End pH 5.07 4.89 4.11 3.61 4.25 3.08
10cm3 of buffer with NaOH:
Few drops 1 cm3 2 cm3 4 cm3 5 cm3 10 cm3
(3)
Start pH 5.08 5.08 5.10 5.09 5.08 5.08
End pH 5.08 5.37 6.92 12.43 12.56 12.95
For 10 cm3 of the buffer solution titrated with HCl:
The initial pH of the buffer was around 5.07-5.08, which is expected for an
ethanoic acid/sodium ethanoate buffer whose pH should be close to the
pKa of ethanoic acid (4.76).
Adding just a few drops (3) of HCl did not change the pH, demonstrating
the buffer's ability to resist small amounts of added acid.
As more HCl was added in 1 cm3 and 2 cm3 quantities, the pH started
decreasing slightly to 4.89 and 4.11 respectively, but not drastically. This
shows the buffer solution neutralizing the added H+ ions.
Larger additions of 4 cm3 and 5 cm3 of HCl caused more significant pH
drops to 3.61 and 4.25, starting to overwhelm the buffer capacity.
Finally, adding 10 cm3 of HCl resulted in a large pH decrease to 3.08,
completely overwhelming and exceeding the buffering range.
For 10 cm3 of buffer titrated with NaOH:
The initial pH was again around 5.08-5.10.
Small additions like a few drops and 1 cm3 of NaOH caused only minor pH
increases to 5.37, showing buffering action.
Adding 2 cm3 of NaOH raised the pH more substantially to 6.92 as the
buffer capacity began getting overwhelmed.
Further additions of 4 cm3 and 5 cm3 of NaOH resulted in large pH
increases to 12.43 and 12.56 respectively, clearly outside the buffering
range.
Finally, 10 cm3 of NaOH completely overwhelmed the buffer, increasing
the pH to 12.95.
In summary, the buffer effectively resisted small amounts of added acid or base,
maintaining the pH around 5. However, as larger volumes of strong acid/base
were added, the buffering capacity was eventually exceeded, resulting in large,
unbuffered pH changes. This demonstrates the buffer action as well as its limits
based on the quantities of acid/base added.
,P3- PRACTICAL 3
1. HCl/NaOH Titration:
Phenolphthalein is an excellent choice for this strong acid-strong base titration
because its colour change pH range of 8.2 - 10.0 aligns perfectly with the vertical
region of the titration curve around pH 7-10.
At the start, the solution is highly acidic due to the hydrochloric acid.
Phenolphthalein remains colourless in this highly acidic pH range below 8.2. As
NaOH is gradually added, it neutralizes the strong acid, consuming H+ ions. The
pH rises steadily until reaching the equivalence point around pH 7.
Near the equivalence point, there is a very sharp rise in pH due to the complete
neutralization of the limiting reagent (acid or base). This vertical jump falls
squarely within the colour change range of phenolphthalein, allowing for a clear
visual indication of the endpoint.
Beyond the equivalence point, the solution becomes basic due to the excess
NaOH, causing phenolphthalein to turn a bright pink/fuchsia color above pH 10.
This distinct color change signals the completion of the titration.
2. HCl/NH3 Titration:
For this weak base titration, methyl red is a judicious choice because its color
change range of pH 4.4 - 6.2 coincides with the buffering region around pH 5-6
on the titration curve.
Initially, the solution is highly acidic due to the hydrochloric acid, and methyl red
appears red. As ammonia (weak base) is gradually added, it starts consuming
H+ ions, raising the pH slowly. The curve exhibits a gradual rise and flattens out
around pH 5-6 due to the buffering action of the resulting NH4+/NH3 system.
During this buffering region, methyl red undergoes a color transition from red to
yellow, providing a clear visual indication of the endpoint. Beyond this point, the
pH rises more steeply as excess ammonia is added, taking the solution into the
basic range where methyl red remains yellow.
3. CH3COOH/NaOH Titration:
Phenolphthalein is again a suitable indicator for this weak acid-strong base
titration. Its colour change range of pH 8.2 - 10.0 aligns well with the steep
vertical region of the titration curve around pH 8-9.
,Initially, the solution is acidic due to the acetic acid, and phenolphthalein
remains colourless. As NaOH is added, it neutralizes the weak acid, consuming
H+ ions and raising the pH gradually. Around the equivalence point, there is a
sharp vertical rise in pH due to the complete neutralization of the limiting
reagent.
This vertical jump falls within the colour change range of phenolphthalein,
allowing for a clear visual indication of the endpoint as the solution turns
pink/fuchsia above pH 8.2. Beyond the equivalence point, the excess NaOH
drives the solution into the basic range, and phenolphthalein remains pink.
4. CH3COOH/NH3 Titration:
For this weak acid-weak base titration, bromocresol green is an appropriate
indicator choice. Its colour change range of pH 3.8 - 5.4 aligns well with the
buffering region around pH 4-6 on the titration curve.
Initially, the solution is acidic due to acetic acid, and bromocresol green appears
yellow. As ammonia is added, it starts neutralizing the weak acid, consuming H+
ions and raising the pH gradually. The curve exhibits a buffering region around
pH 4-6 due to the coexistence of NH4+ and CH3COO- ions.
During this buffering region, bromocresol green undergoes a colour transition
from yellow to blue, providing a clear visual indication of the endpoint. Beyond
this point, the pH rises more steeply as excess ammonia is added, and
bromocresol green remains blue in the basic range.
M2
1. Prepare a buffer solution by partially neutralizing a weak acid (e.g.
ethanoic acid) with its conjugate base (e.g. sodium ethanoate).
2. Measure the pH of the buffer solution using an accurately calibrated pH
meter.
3. Use the Henderson-Hasselbalch equation to calculate the theoretical pH:
pH = pKa + log([base]/[acid]) Where pKa is the negative logarithm of the
acid dissociation constant Ka.
4. Compare the experimental pH to the theoretical value calculated from the
Henderson-Hasselbalch equation. A good buffer should have pH close to
the pKa value.
5. Test buffer action by adding small amounts of strong acid and base
separately to the buffer solution and measuring the pH change. A good
buffer resists large pH changes upon addition of small amounts of H+ or
OH- ions.
, 6. Determine the buffering capacity by continuing to add acid/base and
noting the amounts required to significantly change the pH out of the
buffering range.
Explaining Buffer Action in Blood:
Blood has a pH of around 7.4 which is maintained by a buffering system
based on the carbonic acid/bicarbonate ion equilibrium: H2CO3 ⇌ H+ +
HCO3-
Carbon dioxide produced from cellular respiration dissolves in plasma to
form carbonic acid, providing H+ ions.
Bicarbonate ions (HCO3-) act as the conjugate base to absorb/release H+
and maintain pH.
The Henderson-Hasselbalch equation applies with pKa of 6.1 for this
system.
Blood buffers changes from metabolic acids (e.g. lactic acid from muscles)
or changes in CO2 levels from respiration.
Buffering in blood is crucial as enzymes and other biomolecules are pH
sensitive. Large pH changes can disrupt biological processes.
Excess H+ is removed by breathing out CO2. Kidneys also regulate
bicarbonate levels longer-term to restore the buffering capacity.
So in summary, preparing and testing buffer solutions, applying the Henderson-
Hasselbalch equation, and understanding the carbonic acid-bicarbonate buffering
system in blood would effectively assess and explain buffer action.
M3
1. Titration of 25 cm3 of hydrochloric acid with sodium hydroxide:
This is a strong acid-strong base titration
The titration curve shows a very sharp vertical rise around pH 7 at the
equivalence point
To best indicate the endpoint, the indicator should undergo color change in
the pH range right around 7
Phenolphthalein (color change range 8.2 - 10.0) would be a good choice,
as its range overlaps the vertical portion
Justification: Phenolphthalein is ideal as its color transition from colorless
to pink occurs precisely when the solution changes from acidic to basic
around pH 8-9, allowing clear visualization of the sudden pH change.
2. Titration of 25 cm3 of hydrochloric acid with ammonia solution:
This is a strong acid-weak base titration
, The titration curve has a buffering region around pH 5-6 before showing
the vertical rise
An indicator with color change in the pH 4-7 range would be most suitable
Methyl red (pH range 4.2 - 6.3) or bromocresol green (3.8 - 5.4) would be
appropriate choices
Justification: Methyl red's transition from red to yellow occurs right in the
buffering region where the rapid pH change happens in this titration,
enabling accurate endpoint detection.
3. Titration of 25 cm3 of ethanoic acid with sodium hydroxide:
This is a weak acid-strong base titration
The titration curve has a buffering region around pH 8-9 before the vertical
rise
An indicator with color change in the pH 7-10 range is required
Phenolphthalein (8.2 - 10.0) or thymol blue (8.0 - 9.6) would be good
options
Justification: Phenolphthalein's color change from colorless to pink occurs
exactly in the pH range where the curve has its sharpest vertical portion,
allowing reliable indication of the endpoint.
4. Titration of 25 cm3 of ethanoic acid with ammonia solution:
This is a weak acid-weak base titration
The titration curve exhibits a wide buffering region around pH 4-7
An indicator with color change across this broad range is most suitable
Bromocresol green (3.8 - 5.4) or bromophenol blue (3.0 - 4.6) would be
appropriate choices
Justification: Bromocresol green's transition from yellow to blue occurs
right across the buffering region of this titration, making it an ideal
indicator to signal the gradual pH changes.
DISTINCTION
Titrations Using Acid-Base Indicators
The use of indicators is one of the oldest and most widespread methods for acid-
base titrations. Indicators are weak acids or bases that exhibit different colours
depending on the pH of the solution. When added during a titration, they provide
a visual signal of when the equivalence point is reached by undergoing a distinct
colour change over a narrow pH range.
, One major advantage of indicator titrations is their simplicity and low cost. The
required equipment is very basic - just a burette, conical flask, indicator solution,
and the titrant solution. This makes indicator titrations accessible even in
situations with limited resources. Additionally, the procedure is straightforward,
only requiring the titrant to be slowly added until the colour change is observed.
However, the accuracy of indicator titrations is heavily dependent on choosing
the correct indicator. The assignment brief highlights the importance of proper
indicator selection - it states "You must qualify why you have chosen one
indicator over other possible indicators from your research." Each indicator has a
specific pH range over which its colour transition occurs. This range must overlap
with the expected pH at the equivalence point for the titration to be valid.
Selecting an indicator with a poor pH match can lead to inaccurate or unclear
endpoints.
Even when the right indicator is chosen, there can still be issues impacting
accuracy. The colour change is subjective, relying on the analyst's ability to
discern it. This introduces uncertainty due to human error or differing colour
perception between observers. Furthermore, some indicators provide less distinct
colour transitions compared to others. Certain mixtures of acids/bases may also
interfere with or mask the expected colour change.
Another limitation is that indicator titrations only really detect the approximate
endpoint region rather than pinpointing the precise equivalence point. This can
pose problems when titrating very dilute or poorly buffered solutions where the
pH changes rapidly around the endpoint. Overall, while simple and economical,
relying solely on indicators sacrifices a degree of accuracy compared to other
endpoint detection methods.
Titrations Using a pH Meter
The pH meter offers a more quantitative approach to acid-base titrations
compared to indicators. By directly measuring the pH, it allows the equivalence
point to be accurately identified through examining the titration curve or pH
versus volume data.
One key advantage of pH meters is their high accuracy and precision in dilute
solutions where indicators would struggle. The assignment brief states that for
the titration curves, "As you get closer to the end point, you will need to start to
add amounts of 0.1 cm3 of the alkali so that you can record more readings of pH
to improve the accuracy of your titration curve." These detailed pH readings
generate titration curves with clear equivalence points, allowing the
stoichiometric volume to be reliably determined.
