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CH 125
INORGANIC CHEMISTRY FOR MICROBIOLOGY
Laboratory Manual
2013 Edition
By
Brian Clark, Marco Castillo & Patrick Chan
CENTENNIAL COLLEGE OF APPLIED ARTS &
TECHNOLOGY, SCARBOROUGH, ONTARIO
Preliminary Laboratory Information - CH 125
The following safety information is provided to the student in order to ensure that all students and college staff working in the laboratory are aware of common industrial laboratory safety practices.
Laboratory Safety Wear Requirements:
Each student is required to have the following safety items with them during each laboratory period:
1728. Lab Coat
1729. Safety glasses or goggles
1730. Enclosed shoes
Laboratory Safety Rules:
Each student will be aware of, and abide by the following minimum laboratory safety rules as follows:
1. Each student and teacher must wear safety glasses, lab coat and closed-toe shoes.
2. Do not wear contact lenses in the laboratory. Do not eat, drink or smoke in a laboratory.
3. Know the location, use and operation of the eye-wash fountain, emergency shower, fire extinguisher, fire-blanket, emergency phone and laboratory exits.
4. In case of an accident or cut, immediately notify instructor. Do not leave the laboratory to seek help alone.
5. Clean up all spills at your work area immediately. If a chemical spills on your skin, immediately wash the exposed area under cold water for a minimum of ten minutes. If you have been exposed to a large body spill, quickly proceed to the emergency shower. Remove all affected clothing to ensure chemical spill may be wash from skin.
6. Always sweep up broken glassware using a brush and pan. Keep broken glass in a separate container.
7. Assume all chemicals are hazardous unless instructed otherwise. Never taste a chemical.
8. Do not rub your eyes, mouth or nose unless you know your hands are clean.
9. Corrosive chemicals and chemicals with unpleasant odours are to be kept in the fume hood.
10. Do not heat flammable liquids with an open flame. Always use a hot plate with a hot water bath to heat these liquids.
11. In case of fire, notify all laboratory staff and students immediately by shouting “FIRE”. Turn off all open flames and electrical equipment. Calmly and quickly leave the laboratory. DO NOT FIGHT THE FIRE.
12. If someone is on fire, put the person under the emergency shower, or use the “drop and roll” technique using the fire-blanket.
13. Never point the open end of a test-tube being heated at yourself or another person. The liquid contents may erupt.
14. When inserting glass tubing or thermometers into stoppers or tubing, wet all components with water or lubricate with glycerol. Wrap glass tubing or thermometers in a towel. Gripping the glassware as close as possible to the insertion point, gently twist and ease the glassware into the stopper or tubing.
15. Hot glass and hot “hot-plates” look the same when they are hot or cold. Always approach these items with caution. Place your hand near them to assess whether they are hot before picking them up.
16. Never work alone in a laboratory.
Handling Chemicals and Glassware:
1. Always check the label on the bottle before you use a chemical.
2. To prevent contamination of reagents and chemicals, never insert pipettes or scoopulas into stock bottles. The proper technique is to remove a portion of the stock bottle contents into another piece of clean glassware from which you can obtain your sample. Take only what you need. Do not return any unused chemicals into a stock bottle.
3. Label all glassware containing chemicals with the chemical name, date and your name.
4. Do not pipette a liquid by mouth.
5. Always wash your hands before leaving a laboratory.
6. Do not weigh chemicals directly on a balance. Always use a weighing boat, weighing paper or container.
7. Do not place any chemicals down the sink unless notifies by the instructor. All hazardous chemicals and organic solvents are to be disposed of in their proper disposal containers.
8. Do not assume that any glassware provided is clean. Always wash glassware prior to an experiment and following the experiment.
9. Clean and return all glassware to its proper location before the end of the lab. Do not leave glassware near a sink, fumehood or bench.
Preparation for Laboratory Work:
1. Read the experiment and complete a hand written pre-lab before you come to the laboratory.
2. Arrive to the laboratory on time. All labs will be proceeded by a small lecture regarding the nature of the laboratory, chemical safety and disposal issues. Any student missing a safety precaution lecture will not be allowed to do the lab for safety reasons. Extraordinary circumstances will be accepted.
3. Work independently unless otherwise instructed.
4. Laboratory reports will be submitted at the beginning of the next scheduled laboratory period unless otherwise instructed. There will be a late penalty issues to reports not submitted at this time.
5. There are no makeup labs. Any labs missed will be given a grade of “0" unless accompanied by documented reasons for absences.
CH 125 LABORATORY REPORT GUIDELINES
General Guidelines:
All final work must be done in ink, and should be neat, legible and well organized. Reports should be stapled or preferably be presented in a duo-tang or similar folder (but not a binder). Pre-lab work must be hand written and completed and initialed before the experiment is performed.
Pre-Lab Report will include:
Your Name:
Date Experiment was Performed:
Experiment Number and Title:
Objective:
Safety Precautions:
Procedure: (this must be concise)
Final Report: (To the above information, add the following:)
Observations: (include all data, tables, graphs and all other important observations made during the experiment)
Discussion: (include sources of error, where appropriate)
Conclusion:
Post-Lab Assignment: (answer all questions)
* Please remember that quality is valued over quantity. Be concise and to the point when presenting your report.
* Total value of laboratory portion of this course is 40% of your final grade. * You must obtain a passing grade in both the laboratory and theory sections of the course to pass.
CH 125 LABORATORY SCHEDULE
WINTER 2011
Week
Date
Experiment/Activity
1
7th Jan.
Lab Introduction
2
14th Jan.
Experiment #1 – Cation Analysis
3
21st Jan.
Experiment #2 – Anion Analysis
4
28th Jan.
Experiment #3 – Solubility Determination
5
4th Feb.
Tutorial – Ksp Problem Set
6
11th Feb.
Experiment #4 – pH Indicators and their Inflection Ranges
7
18th Feb.
FAMILY DAY
8
25th Feb.
Experiment #5 – Buffer Systems and Buffer Capacities
4th Mar.
STUDY BREAK
9
11th Mar.
Tutorial – Acid-Base Problem Set
10
18th Mar.
Experiment #6 – Inorganic Redox Reactions
11
25th Mar.
Experiment #7 – Microbiological pH Analysis
12
1st Apr
No LAB
13
8th Apr
Experiment #8 – Microbiological Redox Analysis
Experiment #7 –Follow up
14
15th Apr
Experiment #9 – Microbiological pH & Redox Analyses- TSI Test
Experiment #8 – Follow up
15
22nd Apr
Experiment #9 – Follow up
POLICIES
All laboratory reports are due at the beginning of the following lab period unless the instructor indicates otherwise. The lab instructor, at the beginning of each lab, will detail the write-up requirements. All laboratories must be completed before the end of week 14. Late reports will have a penalty of 20% per week.
Attendance in laboratories is mandatory. Lab reports will not be accepted for missed labs. For labs that missed due to illness, please provide a doctor’s note and make arrangements for the missed work with your instructor.
Experiment #1: Cation Analysis
Introduction:
Chemical analyses are categorized into two specific groups, namely qualitative analysis and quantitative analysis. The purpose of a qualitative analysis is to determine the identity or composition of a compound or mixture of compounds, whereas the goal in quantitative analysis is to determine how much of a given substance is present.
Qualitative analysis of an aqueous mixture typically proceeds in three stages:
1) The ions present in solution are separated into broad groups based on their solubility properties.
2) The individual ions within each group are separated by specific extraction techniques.
3) The identity of the specific ions are confirmed by specific chemical tests.
It is important to note that elements are grouped together in a qualitative analysis based on their solubility properties. Although these are called Group I elements, Group II elements, etc., this terminology should not be confused with the Group nomenclature within the periodic table, which is based on electronic configuration.
In cation analysis, typical qualitative analysis separates the ions present in a solution into six groups as shown in Table 1. Subsequent separation and identification is made possible by careful manipulation of the following:
i) the pH of the solution, ii) the known solubilities of different compounds of the ions at different conditions, iii) specific tests that produce a coloured compound with a particular ion, where the colour observed is characteristic of that ion only.
When conducting a complex qualitative analysis, the order of addition of reagents is important. Analysts commonly proceed with separating and testing for Group I ions, followed by Group II, Group III, and so forth. When conducting this type of analysis, one must ensure that all reactions have proceeded to completion, and that any preceding Group ions remaining in a solution are at such a low concentration that they will not interfere with subsequent tests. For example, a white precipitate formed in Group V cannot be AgCl, since all the Ag+ ions were precipitated in the Group I analysis.
Table 1: Qualitative Cation Analysis Chart
Group
Ions that can be precipitated
Precipitating reagent
Nature of precipitate
I
Ag+, Pb2+, Hg22+
hydrochloric acid
chlorides
II
(A) Hg2+, Cu2+, Cd2+, Bi3+, (Pb2+)
(B) As3+, Sb3+, Sn2+, Sn4+
hydrogen sulphide in presence of H+
hydrogen sulphide in presence of H+
sulphides - insoluble in ammonium sulphide
sulphides - soluble in ammonium sulphide
III
Fe3+, Al3+, Cr3+
ammonium hydroxide in presence of ammonium chloride
hydroxides
IV
Zn2+, Mn2+, Ni2+, Co2+
hydrogen sulphide in presence of OH-
sulphides
V
Ba2+, Sr2+, Ca2+, Mg2+
ammonium carbonate
carbonates
VI
Na+, K+, NH4+
specific tests required
------------
Experiment #1:
Separation & Identification of Group I Cations (Silver Group)
1. Add 5 drops of Group I stock solution to a glass test tube. 2. Add 5 drops distilled water to test tube. 3. Add 2 drops 6M HCl and mix. Insoluble chloride precipitates of silver, mercury (I) and lead will develop. 4. Wash sides of test tube with distilled water until precipitate is below level of solution. 5. Centrifuge test tube (be sure to balance test tube in centrifuge). 6. Test for complete precipitation by adding 1 drop of 6M HCl to the clear supernatant. If a precipitate forms, repeat step #5. 7. Place test tube in a cold water bath to chill sample. Then pour the clear solution above the precipitate out into another test tube. This process is known as decant and the clear solution collected is known as decantate. 8. Wash the precipitate from step #7 by adding 5 drops of distilled water and mix. Centrifuge sample, chill and decant. Put the decantate into the same test tube containing the decantate collected in step #7. The decantate may contain cations from the other groups and should be saved only if other analyses are to be conducted. In this case the decantate can be discarded. Save precipitate since it contains AgCl, PbCl2 and Hg2Cl2.
Separation of Lead from Silver and Mercury(I) & Confirmation of Lead:
To separate lead chloride from silver and mercury(I) chloride precipitates, analysts make use of the fact that PbCl2 is soluble in hot water while AgCl and Hg2Cl2 are not. After separation, the presence of lead ions are confirmed by adding potassium chromate to form the bright yellow insoluble lead chromate precipitate, or by adding sulphuric acid to form a white insoluble lead sulphate precipitate.
9. Add 10 to 15 drop of distilled water to the precipitate saved in step #8 and mix.
10. Place test tube in a hot water bath and allow the test tube solution to become hot. Stir solution frequently. (This is done to dissolve PbCl2.)
11. Centrifuge solution at once and decant as soon as possible. The decantate contains the soluble PbCl2, and the precipitate contains the insoluble AgCl and Hg2Cl2. Save both the decantate and the precipitate.
12. To confirm the presence of Pb2+, cool the decantate and divide into two equal portions.
To one portion, add 1 drop of 0.2M potassium chromate. If the bright yellow PbCrO4 precipitate forms, this confirms the presence of Pb2+.
To the second portion, add 1 drop of 2M sulphuric acid and wait (this reaction may be slow). If the white PbSO4 precipitate forms, this confirms the presence of Pb2+.
Separation and Detection of Silver and Mercury(I)
The separation of insoluble silver and mercury(I) chlorides is based on the fact that AgCl is soluble in ammonium hydroxide while Hg2Cl2 is not. Hg2Cl2 reacts with ammonium hydroxide to form the combined products of Hg and HgNH2Cl , both of which are insoluble. (Note: the presence of the greyish Hg and HgNH2Cl precipitate confirms the presence of Hg22+.)
The presence of Ag+ is confirmed by acidifying the solution of Ag+ in NH4OH and observing the precipitation of AgCl.
13. Add 4 drops of 15M NH4OH to the precipitate in step #11, mix well, centrifuge and decant. The decantate contains the soluble Ag+ ions. The presence of a grey or black-gray precipitate confirms the presence of Hg22+.
14. To the decantate in step #13, add 16M HNO3 dropwise, until the solution turns slightly acidic (confirm with litmus paper). A white precipitate confirms the presence of silver.
15. REPEAT STEPS 1 TO 15, SUBSTITUTING YOUR UNKNOWN FOR THE STOCK SOLUTION IN STEP #1. (Use only five drops of your unknown just like the stock solution. Your unknown may contain 1, 2 or 3 cations.))
Analysis Scheme for Group I Cations (Silver Group)
Ag+, Pb2+, Hg22+
(+ 6M HCl)
Decantate
ppt: AgCl PbCl2 Hg2Cl2 Cations of Groups II through VI colour: white white white
(+hot water)
Decantate
ppt: AgCl Hg2Cl2 Pb2+ colour: white white
(+H2SO4) (+K2CrO4)
ppt: PbSO4 PbCrO4 colour: white yellow Pb2+ present
(+NH4OH)
Decantate ppt: Hg + HgNH2Cl Ag(NH3)2Cl colour: grey
Hg22+ present
(+HNO3)
ppt: AgCl colour: white
Ag+ present
Analysis Scheme for Group I Cations (Known Solution)
Group I Cations
______, ______, ______
(+ 6M HCl)
Decantate
ppt: Yes or No ppt may contain: ______ ______ ______ Colour: ______ colour: ______ ______ ______
(+hot water)
Decantate ppt: Yes or No ppt may contain: ______ ______ Colour: ______ colour: ______ ______ May contain:______
(+H2SO4) (+K2CrO4)
ppt: ______ ______ colour: ______ ______
______ present
(+NH4OH)
Decantate ppt: ______ Colour:______ colour: ______ May contain: ______
______ present (+HNO3)
ppt: ______ colour: ______ ______ present
Analysis Scheme for Group I Cations (Unknown Solution #______)
Group I Cations
______, ______, ______
(+ 6M HCl)
Decantate
ppt: Yes or No ppt may contain: ______ ______ ______ Colour: ______ colour: ______ ______ ______
(+hot water)
Decantate ppt: Yes or No ppt may contain: ______ ______ Colour: ______ colour: ______ ______ May contain:______
(+H2SO4) (+K2CrO4)
ppt: ______ ______ colour: ______ ______
______ present
(+NH4OH)
Decantate ppt: ______ Colour:______ colour: ______ May contain: ______
______ present
(+HNO3)
ppt: ______ colour: ______
______ present
Experiment #2: Anion Analysis
Introduction:
As with cation analyses, anion analyses are also categorized into two specific groups previously discussed, namely qualitative analysis and quantitative analysis. However, each anion is tested for individually based on a sample taken from the original unknown solution.
For this experiment, specific anion tests for the qualitative determination of carbonates, chlorides, sulphates, nitrates and phosphates are provided. The specific chemistry of each test is as follows:
Carbonates: CO32- (aq) + 2H+ (aq) CO2 (g) + H2O
Chlorides: Cl- (aq) +Ag+ (aq) AgCl (s) (white ppt)
Sulphates: SO42- (aq) + BaCl2 (aq) BaSO4 (s) + 2Cl- (aq) (white ppt)
Nitrates: 8Fe2+ (aq) + 2NO3- (aq) + 8H+ (aq) 2Fe(NO)2+ + 6Fe3+ (aq) + 4H2O
(Note: Reduction of NO3- to NO by iron(II) results in formation of brown Fe(NO)2+ complex)
Phosphates: PO43- (aq) + 12MoO42- (aq) + 24H+ (aq) + 3 NH4+ (aq) (NH4)3 PO4.12MoO4 (s) + 12 H2O (yellow ppt)
Procedures:
Carbonates: Place about 20 drops volume of aqueous carbonate solution or solid mixture containing carbonates into a test tube. Add 3 to 4 drops of 6M HCl and observe. If a rapid effervescence occurs with a solid, or if bubbles form in an aqueous solution, this may indicate the formation of CO2 gas, indicating the presence of carbonates in the unknown. (If one wants to confirm the gas produced is carbon dioxide, allow the gas to come in contact with a saturated solution of Ba(OH)2 or Ca(OH)2. The formation of a white precipitate will confirm the gas produced is carbon dioxide.)
Chlorides: Place about 20 drops of solution in a test tube, then add 3 to 4 drops of 3M HNO3 until the solution is slightly acidic. Add 1 or 2 drops of 0.2M AgNO3 and mix. If a white precipitate forms, which is subsequently soluble in ammonium hydroxide, then Cl- ions were present in the original sample.
Sulphates: Place 20 drops of solution in a test tube, then add 3 to 4drops of 3M HNO3 until the solution is slightly acidic. Add 2 to 3 drops of 0.2 M BaCl2 and mix. If a white precipitate forms, then SO42- ions were present in the original sample.
Nitrates: Place 20 drops of sample into a test tube, then add a few crystals of solid ferrous sulphate or a few drops of aqueous ferrous sulphate solution. Carefully add about 2 mL of concentrated sulphuric acid by pouring it down the wall of the test tube while test tube is at a 45o angle. When this step has been completed, there will be two layers of solution in the test tube. The appearance of a brown ring at the interface between the two solutions indicates the presence of the NO3- ion in the original sample.
Phosphates: Place 20 drops of sample in a test tube, then add 4 to 5 drops of 3M nitric acid and 3 to 4 drops of 0.2 M ammonium molybdate solution. Place test tube in a hot water bath and observe. The formation of a slowly forming yellow-coloured solution or precipitate indicates the presence of the PO43- ion in the sample.
REPEAT ALL ANALYSES FOR THE UNKNOWN PROVIDED BY INSTRUCTOR.
Experiment #3: Solubility Determination
Introduction:
Many inorganic compounds are only slightly soluble in water. When these compounds are dissolved in water, the solution will become saturated with the cations and anions present in the compound.
i.e. CuCl(s) Cu+(aq) + Cl-(aq)
At this point, we can state that the ions in solution are in equilibrium with the solid, and the ion concentrations present in the water are at their greatest concentration. Addition of any more cations or anions will result in the precipitation of the slightly soluble solid.
The equilibrium solubility concentrations of both the cations and anion in solution may be expressed using a Solubility Product Constant. This constant may be determined for any slightly soluble compound having the form AxBy as follows:
For the solubility reaction: AxBy(s) x A+(aq) + y B-(aq)
Ksp = [A+(aq) ]x [B-(aq) ]y
where [ ] refers to the ion concentrations in units of molarity (mol/L).
The value of the Ksp constant is highly dependent on temperature, and is always reported at a specific temperature. (i.e. Ksp of CuCl = 3.2x10-7 @ 25oC). In addition, the solubility product constant indicated that it is the “product” of the cation and anion concentrations that determine if precipitation will occur, so the analyst must guard against looking at one ion independently of the other.
In this lab, the students will familiarize themselves with the concept of equilibrium solubility by determining the solubility of PbCl2 at various temperatures. From these data, the student will estimate the solubility product of PbCl2 and compare it to the literaure value.
Ksp of PbCl2 = [Pb2+][Cl-]2
[Pb2+] = molar solubility of PbCl2
[Cl-] = 2 x molar solubility of PbCl2
Procedure:
1. Prepare a boiling bath using 300 to 400 mL of tap water in a 600 mL beaker.
2. Using weighing paper, carefully weigh PbCl2 (to the nearest +/- 1 mg) in the amounts shown in the table below, and transfer into 5 marked, dry, large test tubes.
*Note: It is important that the test tubes are dry and you transfer the PbCl2 into the test tubes first before you add the water.
Tube # Sample weight Solubility Molar Solubility Temp. (g) (g/100g) (mol/L) (oC)
#1 0.500
#2 0.450
#3 0.400
#4 0.350
#5 0.300
3. Add 25.0 mL distilled water to each of the samples in the tubes, mix well, and mark the liquid level on the test tube. Place all test tubes in the boiling water bath ensuring that the liquid level in the test tube is below the liquid level of the boiling water bath.
4. When any solution has completely dissolved (this may require some mixing using a stirring rod), remove the numbered test tube from the boiling water bath, quickly check to make sure the liquid level in the test tube has not changed. If the level has dropped, bring the level back to the original level with distilled water. Insert a thermometer into the solution. and measure the temperature at which the first crystals begin to precipitate. Record the temperature.
5. Repeat step #4 for all solutions. (Note that Tube solution #5 will need to be cooled under cold running water.)
Analysis of Data:
Once all 5 crystallisation temperatures have been found, tabulate your results using the Table shown above . This includes the conversion of your solubility data into g PbCl2/100 g solution (assuming a solution density of 1.00 g/mL). Plot Molar Solubility of PbCl2 sol. vs. Temperature in oC .
Determine the molar solubility of PbCl2 @ 25oC from the graph. Calculate the Ksp (solubility product) at 25oC from the molar solubility and compare this value to the literature values which is 1.6 x 10-5.
Experiment #4: pH Indicators and their Inflection Ranges
Introduction:
In any biological or microbiological system, the pH of the solution is an important factor as to whether the species will be able to survive in its environment. (For example, when discharging water to Lake Ontario, the pH must be between 6.5 and 8.5, otherwise the fish population may be adversely affected.) The pH of any solution is defined by the mathematical equation pH = -log[H+]. To measure the pH of a solution, an analyst may use either a pH meter or a variety of pH indicators.
This lab introduces the concept of pH and indicators commonly used in chemistry and microbiology. Students will determine inflection ranges of known and unknown indicators, as well as using pH meters. They will compare their results with data reported in the literature.
Procedure:
Using the 4 common pH indicators, litmus, brilliant green, phenolphthalein and thymol blue, determine the inflection range and colour changes observed for each of the indicators. For consistency, when you record your results, always write down the colour change from acid to base and NOT in reverse. 1. Pre-test: Start with 2 test tubes, each 1/3 full with distilled water and add 1 drop of one indicator to each tube. Add 1 drop of 0.1 M HCl to one test tube and 1 drop of 0.1 M NaOH to the other tube and observe the colour changes, if any. If no colour changes occur, add 1drop of 1 M HCl and 1 M NaOH respectively to the same tubes and observe colour changes. Repeat for the other indicators.
2. Place 50 mL of distilled water in a 100 mL beaker. Using a stirring rod for mixing, add 8-10 drops of one indicator and record the solution colour. Depending upon your results in step #1, add 1 drop of either 0.1 M HCl or 0.1M NaOH and observe any colour change, if any. If no colour change is apparent, continue to add more drops of acid or base until the colour changes, but keep track of the number of drops added. (Note: If the pre-test indicated the use of a strong acid or base, use this instead.) Discard the solution and repeat for the other indicators.
3. Test: Once the approximate number of drop required is known, again place 50 mL of distilled water with 8-10 drops of indicator in a 100 mL beaker on top of a magnetic stir plate. Put a stir bar and insert a standardized pH probe into the beaker. The starting pH should be about 6. Turn stirrer on by using the left knob. (Do not turn on heat.) Add the appropriate number of drops of acid or base required for colour changes as determined in the previous step. Record the pH for the colour change. For best results, use fractions of drops. Always use the squeeze bottle or swirl the contents to wash down reagent drops adhering to the beaker walls. Obtain the precise upper and lower pH values at which the indicators no longer change colour. (If you overshot the colour change, you can go back by using acid or base.) Repeat for the other indicators.
Note: Before using the pH meter, standardize it. Set it to 25oC, then use a buffer solution with a constant pH, e.g. pH = 7.00. Rinse the electrode after switching to Stand-by. Always store the electrode in buffer solution. To check the pH of a solution, immerse the electrode into the solution after rinsing the probe with distilled water. Swirl solution for approximately 20 to 30 seconds and then take the pH reading.
Note: Thymol Blue has 2 inflection ranges.
4. Record colour changes and pH range of the inflection point.
5. Determine the identity of an unknown indicator by repeating steps #1 to 3 as previously discussed. Base the identity on both colour changes and pH inflection point data. Record
LIST OF COMMONLY USED pH INDICATORS
Name pH Range Colour (acidic to basic)
Crystal Violet 0.0 - 1.8 yellow to blue
* Brilliant Green yellow to green
Malachite Green 1.0 - 1.8 yellow green to blue green
*Thymol Blue red to yellow
Crystal Violet 1.6 - 2.5 green to violet
Methyl Violet 1.8 - 2.5 violet to blue
Methyl Orange 3.2 - 4.4 pink to yellow
Bromophenol Blue 3.0 - 4.6 slight yellow to blue
Congo Red 3.0 - 5.0 violet to red/orange
Bromocresol Green 3.8 - 5.4 yellow to blue
Methyl Red 4.8 - 6.0 pink to yellow
Mod. Methyl Red 4.3 - 5.8 violet to blue/green
* Litmus red to blue
Bromocresol Purple 5.2 - 6.8 slight yellow to purple
Bromothymol Blue 6.0 - 7.6 yellow/orange-green-blue
Neutral Red 6.8 - 8.0 purple to slight yellow
Phenol Red 6.6 - 8.0 violet to orange/red
Thymol Blue yellow to blue m-Nitrophenol 6.8 - 8.6 colourless to yellow
* Phenolphthalein colourless to pink
Thymolphthalein 9.4 - 10.6 colourless to blue
Alizarin Yellow 10.0-12.0 clear to deep yellow
Known indicators used in the experiment. pH inflection range not reported.
Experiment #5: Buffer Systems and Buffer Capacities
Introduction:
A buffer solution is an equilibrium system which, within limits, resists changes to pH upon the addition of strong acids or strong bases. Buffer systems are made when a weak acid or weak base is mixed with its corresponding salt. Buffers are commonly incorporated into microbiological growth media to ensure the pH range of the solution remains within the acceptable pH range for biological growth.
The capacity of a buffer is the amount of acid or base required to change the pH of 1 litre of buffer by one pH unit. It is expressed in units of moles of acid or base per litre buffer. Depending upon the pH, acidic buffers are titrated with a base and basic buffers are titrated with acid.
This lab familiarizes the student with the concept of commonly used buffer systems. Students will produce buffers for specified pH and perform buffering capacity measurements and compare their results with literature data.
Procedure:
1. Prepare a Butterworth Buffer from the table below according to the pH specified by the instructor. Using a volumetric cylinder and a Mohr pipette, the appropriate volumes of the two buffer components are mixed and measured using the previously calibrated pH meter. The measured pH should agree within +/- 0.2 pH units. Record results.
Butterworth Buffer Table for pH 5.80 to pH 8.00 Buffers pH x pH x
5.80 3.6 7.00 29.1
5.90 4.6 7.10 32.1
6.00 5.6 7.20 34.7
6.10 6.8 7.30 37.0
6.20 8.1 7.40 39.1
6.30 9.7 7.50 41.1
6.40 11.6 7.60 42.8
6.50 13.9 7.70 44.2
6.60 16.4 7.80 45.3
6.70 19.3 7.90 46.1
6.80 22.4 8.00 46.7
6.90 25.9
Note: Buffer solution made using 50 mL of 0.1M KH2PO4 and x mL of 0.1M NaOH. Final volume of all buffer mixtures is 100 mL.
Since the given NaOH solution is not exactly 0.1 M it is necessary to calculate and use the equivalent volume.
2. Once the buffer has been satisfactorily prepared and shown to the instructor, the student proceeds with the buffer capacity measurement. Using a pH meter, stirrer bar and magnetic stirrer plate, and a burette mounted on a stand, fill the burette to exactly zero mL with the given standardized 0.1M HCl. The lower meniscus is to be taken as your liquid level. Use 50 mL of your buffer solution, as measured with a graduated cylinder. The remaining 50 mL of your buffer solution is use for step #4.
3. Add reagent slowly, drop by drop,(one drop is approximately 0.05 mL) from your burette until a pH change of 0.1 units has been observed. Read your reagent consumption from your burette and record it. Continue process in intervals of 0.1 pH units, until the buffer system’s pH has changed by 1.1 units.
4. Repeat steps #2 and 3 with the remaining 50 mL of your buffer solution, but change the reagent in the burette to the given standardized 0.1M NaOH.
5. From the volume of HCl and NaOH added find the # of mol of H+ and OH- added for each data point. Plot pH vs # of mol of H+ and # of mol of OH- as shown in the figure below.
6. From the plot obtain the # of mol of H+ and OH- required to change the pH by 1.0 pH unit for 50 mL of your buffer solution. Calculate your nominal buffer capacity for a change in 1.0 pH units in # of mol of H+ (acid) and OH- (base) per litre of buffer solution.
pH
# mol of H+ added 0 # of mol of OH- added
Plot of pH vs mol of HCl and mol of NaOH.
Experiment #6: Inorganic Redox Reactions
Introduction:
Biochemical redox (reduction-oxidation) reactions are commonly used in microbiology to identify microorganisms. These reactions are often based on metabolic products formed by specific microorganisms, which can then be subsequently identified. For example, nitrate reducing bacteria may convert the nitrate ions (NO3-) present in water to nitrite ions (NO2-) under anaerobic conditions. To identify the nitrate reducing bacteria, an analyst would inoculate a nitrate broth with bacteria and subsequently monitor the solution for the formation of nitrites.
At other times, specific compounds are added to a growth medium which can only be utilized by certain micro-organisms which are often growth inhibiting (toxic) to other microorganisms. One common example of these types of compounds is oxygen which separates aerobes from the anaerobes.
The purpose of this laboratory is to familiarize the student with the concept of redox reactions, and introduce the student to typical redox indicators used in microbiological analyses.
Procedure:
1. Methylene Blue Redox Test
A. Fill a test tube 1/4 full with distilled water. Add three to five drops of methylene blue, three drops of 6 M hydrochloric acid and a pinch of Zn dust. The evolution of hydrogen gas reduces the methylene blue dye to a colourless reduction product.
Chemical Reaction: 2HCl(aq) + Zn(s) H2(g) + ZnCl2(aq)
Redox Reactions: Zn Zn2+ + 2e-
2H+ + 2e- H2 methylene blue (oxidized) + 2e- methylene blue (reduced)
(blue) (colourless)
B. Decant the clear solution from the zinc dust by pouring it slowly in another test tube. Add 2 drops of 3% hydrogen peroxide (H2O2) solution. Record your results.
Reaction: methylene blue (reduced) + H2O2 + 2H+ methylene blue (oxidized)+
2H2O
2. Nitrate Reduction Test:
Using 4 test tubes, each filled to ¼ full with distilled water, do the following tests and record your observations.
Tube #1: 3 drops Nitrate solution(1%) + 3 drops Nitrite ID agent + 1 drop 5M HCl
KNO3 + nitrite indicator + HCl no reaction
Tube #2: 3 drops Nitrate solution(1%) + 3 drops Nitrite ID agent + 1 drop 5M HCl + Zn
KNO3 + 2HCl + Zn ZnCl2 + KNO2 + H2O
KNO2 + nitrite indicator azodye (pink/red)
Tube #3: 3 drops Nitrite solution(1%) +3 drops Nitrite ID agent + 1 drop 5M HCl
KNO2 + nitrite indicator azodye (pink/red)
Tube #4: 3 drops Nitrite solution(1%) +3 drops Nitrite ID agent + 3 drop 5M HCl + Zn
NaNO2 + 3Zn + 8HCl NH4Cl + 3ZnCl2 + 2H2O + NaCl
The pink-red colour observed in tubes #2 and 3 is based on the formation of an azodye, which is formed when nitrite is present.
Note: Heating the test tubes may be necessary to speed up reaction. If required, place test tubes in a hot water bath.
3. Hydrogen sulfide formation from thiosulphate:
Into each of 3 test tubes, each filled to ¼ full with distilled water, add a few drops of sodium thiosulphate solution(5%) + ferrous sulfate solution(5%). Add the following as shown, and record your observations.
Tube #1: Add a pinch of zinc dust plus 1 drop dilute HCl.
Na2S2O3 + 3Zn + 8HCl H2S + 3ZnCl2 + S0 + 3H2O + 2 NaCl
Redox Reactions: S2O32- + 6H+ + 6e- S0 + S2- + 3H2O
Zn Zn2+ + 2e-
Test for evolution of H2S by gently smelling the contents of the test tube. Hydrogen sulphide has a characteristic “rotten egg” smell.
Proof for H2S formation: FeSO4 + H2S + 2NH3 FeS (black ppt) + (NH4)2SO4 Tube #2: Add only a pinch of zinc dust:
Na2S2O3 + Zn no reaction
Tube #3: Add only 1 drop dilute HCl:
Na2S2O3 + HCl S0 + H2SO3 + 2NaCl
S2O32- S0 (yellow/white ppt) + SO32-
4. Testing Redox Agents with Starch and Iodine/Iodide: Fill 2 test tubes to ¼ full with room temperature distilled water and 1 test tube with cold distilled water. Add 2 drops of 10% acetic acid to each of the test tubes. Do the following tests and record your results.
Tube #1: Add 3 drops of iodine (KI-I2) solution plus 3 drops of starch solution to one of the test tubes with room temperature water. Note the purple to blue colour formed. The starch is being used to indicate the presence of iodine.
Tube #2: Add 3 drop of potassium iodide (KI) solution plus 3 drops of starch solution to the test tube with cold distilled water. Record initial observations. Shake the tube vigorously for 1-2 minutes and observe again. You should see a slight blue colouration after a while, based on atmospheric oxidation of iodide to elemental iodine. (This procedure works better with cold water.)
2H+ + ½O2 + 2I- I2 + H2O
Redox Reactions: 2H+ + ½O2 + 2e- H2O
2I- I2 + 2e-
Tube #3: Add 3 drops of iodine (KI-I2) solution plus 3 drops of starch to the test tube with room temperature distilled water. Observe, then add several drops of thiosulphate (Na2S2O3) solution and observe again.
2Na2S2O3 + KI-I2 + 2HAc Na2S4O6 + KI + 2HI + 2NaAc
As in the previous cases, starch is used as an indicator for the presence of iodine. When sodium thiosulphate is added, the iodine is reduced to iodide, and will not test positive with the starch test. Note: This reaction is used in quantitative analytical iodometric titrations (i.e sulphite tests when determining presence of sulphate and sulphite reducing bacteria).
Experiment #7: Microbiological pH Analysis
Introduction:
In many microbiological systems, bacteria will ferment a carbohydrate broth to produce various acidic byproducts. Depending upon the carbohydrate used, the initial pH of the broth, presence of various organic and/or inorganic compounds, the type of bacteria, etc., the identification of the microorganism may be determined. For example, in the mannitol salt medium, only bacteria that ferment the sugar mannitol in a concentrated salt broth will survive (typically staphylococci). By using a pH indicator to detect the presence of acidic byproducts, a microbiologist may identify the presence of staphylococci using this media.
In this experiment, the student will prepare various phenol red carbohydrate media and apply his/her knowledge of pH indicators (Expt. #4) in determining the function of various simple microbiological media.
Procedure:
Each group of students will be assigned a broth to prepare and share with the other groups. Each broth will have a volume of 250 mL and will consist of the following:
Ingredients
Base Broth
Mannitol Broth
Lactose Broth
Maltose Broth
Peptone
2.5g
2.5g
2.5g
2.5g
Beef Extract
0.25g
0.25g
0.25g
0.25g
Mannitol
1.25g
Lactose
1.25g
Maltose
1.25g
Sodium Chloride
1.25g
1.25g
1.25g
1.25g
Phenol Red
2 mL
2 mL
2 mL
2 mL
Water
to 250 mL
to 250 mL
to 250 mL
to 250 mL Gently boil the medium for 10 minutes to sterilize the medium and drive off any dissolved gases. Wait until the medium has cooled down sufficiently adjust the pH of each broth to 7.4 ± 0.2 by adding 0.1 M.NaOH or 0.1 M HCl dropwise. Pour medium into culture vials, and add a Durham tube to each vial to assess if gas formation occurs. Label each vial to indicate the type of broth and give one vial to each of the other groups.
Record initial condition and colour of vial. Record final condition(s) and colour during next laboratory session.
NOTE: Mark your individual group cultures carefully. Each culture will be inoculated with bacteria after being autoclaved. They will be incubated in the microbiology laboratory. Incubated vials will be available to students at the next scheduled laboratory session for observations.
Post-Laboratory Assignment:
1. Based on your knowledge of microorganisms, briefly discuss whether the bacteria used to inoculate each broth were able to ferment the carbohydrates present in the broth. On what basis are these statements made? 2. What is the purpose of each ingredient in the broth mixture assigned to your group?
3. If phenol red was not available for preparing the medium, what suitable pH indicator(s) may be used? How would this impact on the test?
Experiment #8: Microbiological Redox Analysis
Introduction:
Some bacteria, such as Salmonella sp. and Proteus sp., are able to produce hydrogen sulphide (H2S) by reducing inorganic compounds such as thiosulphate (S2O32-), sulphates, (SO42-) , and/or sulphites (SO32-). In each case, the oxidation state of sulphur changes from the +2, +6 and +4 respectively to the -2 oxidation state in the acidic conditions associated with the bacteria. To use this fact as a basis of determination, the SIM medium is used to promote growth of bacterial strains that are capable of reducing thiosulphate to hydrogen sulphide.
In this experiment, the student will prepare a SIM medium, and apply his/her knowledge of redox indicators (Expt. #6) in determining the function of this microbiological media.
Procedure:
Each group of students will prepare 200 mL of SIM medium as follows:
Ingredients
SIM Media
Peptone
6.0g
Beef Extract
0.6g
Ferrous Sulphate
0.04g
Sodium Thiosulphate
0.005g
Agar
0.7g
Water
to 200 mL Gently boil the medium for 10 minutes to sterilize the medium and drive off any dissolved gases. Wait until the medium has cooled down sufficiently adjust the pH of each broth to 7.3 ± 0.2 by adding 0.1 M.NaOH or 0.1 M HCl dropwise. Pour media into four (4) culture vials to about half full and allow the medium to solidify in the vertical position. Label each vial to indicate the name of the group and sample number.
Record initial condition and colour of vial. Record final condition(s) and colour during next laboratory session.
NOTE: Mark your individual group cultures carefully. Each culture will be inoculated (stab) with bacteria after being autoclaved. They will be incubated in the microbiology laboratory. Incubated vials will be available to students at the next scheduled laboratory session for observations.
Post-Laboratory Assignment:
1. Based on your knowledge of microorganisms, briefly discuss whether the bacteria used to inoculate each broth were able to reduce the thiosulphate present in the medium? On what basis are these statements made? 2. What is the purpose of each ingredient in the medium?
Experiment #9: Microbiological pH & Redox Analyses - TSI Test
Introduction:
When provided with a sample of unknown bacteria, an analyst will often rely on test media that provide information on both the types of carbohydrates a bacteria can ferment, as well as the microorganism’s ability to reduce various inorganic compounds, as an aid in the identification process. The Triple Sugar-Iron Agar test (TSI) is a rapid screening test that is able to differentiate and distinguish the members of the Enterobacteriaceae and other intestinal bacilli (intestinal bacteria found in contaminated food and water supplies).
The TSI medium utilizes the three sugars at the following concentration - lactose (1%), sucrose (1%) and glucose (0.1%), as well as sodium thiosulphate to provide differentiation among the various bacteria.
In this experiment, the student will prepare a TSI medium and apply his/her knowledge of pH and redox indicators in determining the function of this microbiological media.
Procedure:
Each group of students will prepare 200 mL of TSI media as follows:
Ingredients
TSI Media
Ingredients
TSI Media
Beef Extract
0.6g
Sucrose
2.0g
Yeast Extract
0.6g
Ferrous Sulphate
0.04g
Peptone
3.0g
Sodium Chloride
1.0g
Proteose Peptone
1.0g
Sodium Thiosulphate
0.06g
Dextrose
0.2g
Agar
2.4g
Lactose
2.0g
Phenol Red
5 mL Gently boil the medium for 10 minutes to sterilize the medium and drive off any dissolved gases. Wait until the medium has cooled down sufficiently adjust the pH of each broth to 7.4 ± 0.2 by adding 0.1 M.NaOH or 0.1 M HCl dropwise. Pour medium into four (4) culture vials to about half full. Label each vial to indicate the name of the group and sample number.
Record initial condition and colour of vial. Record final condition(s) and colour during next laboratory session.
NOTE: Mark your individual group cultures carefully. Each culture will be inoculated (stab and streak) with bacteria after being autoclaved and allowed to solidified in a slant position. They will be incubated in the microbiology laboratory. Incubated vials will be available to students at the next scheduled laboratory session for observations.
Post-Laboratory Assignment:
1. Based on your knowledge of microorganisms, briefly discuss whether the bacteria used to inoculate each broth were able to ferment the carbohydrates and reduce the thiosulphate present in the medium? On what basis are these statements made? 2. What is the purpose of each ingredient in the medium?
3. Originally, this test utilized bromothymol blue (inflection point: 6.0-7.6) as the pH indicator. How would this impact on the test?
4. In the original industrial production of penicillin, the strain Penicillium chrysogenum was introduced into a culture medium consisting of the following:
Grams per Litre Medium
Corn liquor 30.0 g
Lactose 30.0 g
Glucose 5.0 g
NaNO3 3.0 g
MgSO4 0.25 g
ZnSO4 0.044 g
Phenyl acetamide 0.05 g
(penicillin precursor)
CaCO3 3.0 g
NaHCO3 3.0 g
The bacterial strain would utilize the food and nutrients in the media, in the presence of oxygen from air bubbled through the mixture, to produce penicillin. The penicillin produced was subsequently extracted and purified for use.
a) What is the purpose of each of the ingredients present in the medium?
b) Calculate the pH of the medium based on the presence of the carbonate ion in equilibrium with the bicarbonate ion according to the reaction:
HCO3- H+ + CO32- Ka2 = 5.61 x 10-11
Hint: Assume all the sodium bicarbonate in the medium is soluble. Determine the concentration of the carbonate ions based on the solubility product of CaCO3 .
c) If the laboratory did not have ZnSO4 available, could it be replaced with ZnS on an equivalent molar basis? What concerns would you have?
INORGANIC CHEMISTRY FOR MICROBIOLOGY
Laboratory Manual
2013 Edition
By
Brian Clark, Marco Castillo & Patrick Chan
CENTENNIAL COLLEGE OF APPLIED ARTS &
TECHNOLOGY, SCARBOROUGH, ONTARIO
Preliminary Laboratory Information - CH 125
The following safety information is provided to the student in order to ensure that all students and college staff working in the laboratory are aware of common industrial laboratory safety practices.
Laboratory Safety Wear Requirements:
Each student is required to have the following safety items with them during each laboratory period:
1728. Lab Coat
1729. Safety glasses or goggles
1730. Enclosed shoes
Laboratory Safety Rules:
Each student will be aware of, and abide by the following minimum laboratory safety rules as follows:
1. Each student and teacher must wear safety glasses, lab coat and closed-toe shoes.
2. Do not wear contact lenses in the laboratory. Do not eat, drink or smoke in a laboratory.
3. Know the location, use and operation of the eye-wash fountain, emergency shower, fire extinguisher, fire-blanket, emergency phone and laboratory exits.
4. In case of an accident or cut, immediately notify instructor. Do not leave the laboratory to seek help alone.
5. Clean up all spills at your work area immediately. If a chemical spills on your skin, immediately wash the exposed area under cold water for a minimum of ten minutes. If you have been exposed to a large body spill, quickly proceed to the emergency shower. Remove all affected clothing to ensure chemical spill may be wash from skin.
6. Always sweep up broken glassware using a brush and pan. Keep broken glass in a separate container.
7. Assume all chemicals are hazardous unless instructed otherwise. Never taste a chemical.
8. Do not rub your eyes, mouth or nose unless you know your hands are clean.
9. Corrosive chemicals and chemicals with unpleasant odours are to be kept in the fume hood.
10. Do not heat flammable liquids with an open flame. Always use a hot plate with a hot water bath to heat these liquids.
11. In case of fire, notify all laboratory staff and students immediately by shouting “FIRE”. Turn off all open flames and electrical equipment. Calmly and quickly leave the laboratory. DO NOT FIGHT THE FIRE.
12. If someone is on fire, put the person under the emergency shower, or use the “drop and roll” technique using the fire-blanket.
13. Never point the open end of a test-tube being heated at yourself or another person. The liquid contents may erupt.
14. When inserting glass tubing or thermometers into stoppers or tubing, wet all components with water or lubricate with glycerol. Wrap glass tubing or thermometers in a towel. Gripping the glassware as close as possible to the insertion point, gently twist and ease the glassware into the stopper or tubing.
15. Hot glass and hot “hot-plates” look the same when they are hot or cold. Always approach these items with caution. Place your hand near them to assess whether they are hot before picking them up.
16. Never work alone in a laboratory.
Handling Chemicals and Glassware:
1. Always check the label on the bottle before you use a chemical.
2. To prevent contamination of reagents and chemicals, never insert pipettes or scoopulas into stock bottles. The proper technique is to remove a portion of the stock bottle contents into another piece of clean glassware from which you can obtain your sample. Take only what you need. Do not return any unused chemicals into a stock bottle.
3. Label all glassware containing chemicals with the chemical name, date and your name.
4. Do not pipette a liquid by mouth.
5. Always wash your hands before leaving a laboratory.
6. Do not weigh chemicals directly on a balance. Always use a weighing boat, weighing paper or container.
7. Do not place any chemicals down the sink unless notifies by the instructor. All hazardous chemicals and organic solvents are to be disposed of in their proper disposal containers.
8. Do not assume that any glassware provided is clean. Always wash glassware prior to an experiment and following the experiment.
9. Clean and return all glassware to its proper location before the end of the lab. Do not leave glassware near a sink, fumehood or bench.
Preparation for Laboratory Work:
1. Read the experiment and complete a hand written pre-lab before you come to the laboratory.
2. Arrive to the laboratory on time. All labs will be proceeded by a small lecture regarding the nature of the laboratory, chemical safety and disposal issues. Any student missing a safety precaution lecture will not be allowed to do the lab for safety reasons. Extraordinary circumstances will be accepted.
3. Work independently unless otherwise instructed.
4. Laboratory reports will be submitted at the beginning of the next scheduled laboratory period unless otherwise instructed. There will be a late penalty issues to reports not submitted at this time.
5. There are no makeup labs. Any labs missed will be given a grade of “0" unless accompanied by documented reasons for absences.
CH 125 LABORATORY REPORT GUIDELINES
General Guidelines:
All final work must be done in ink, and should be neat, legible and well organized. Reports should be stapled or preferably be presented in a duo-tang or similar folder (but not a binder). Pre-lab work must be hand written and completed and initialed before the experiment is performed.
Pre-Lab Report will include:
Your Name:
Date Experiment was Performed:
Experiment Number and Title:
Objective:
Safety Precautions:
Procedure: (this must be concise)
Final Report: (To the above information, add the following:)
Observations: (include all data, tables, graphs and all other important observations made during the experiment)
Discussion: (include sources of error, where appropriate)
Conclusion:
Post-Lab Assignment: (answer all questions)
* Please remember that quality is valued over quantity. Be concise and to the point when presenting your report.
* Total value of laboratory portion of this course is 40% of your final grade. * You must obtain a passing grade in both the laboratory and theory sections of the course to pass.
CH 125 LABORATORY SCHEDULE
WINTER 2011
Week
Date
Experiment/Activity
1
7th Jan.
Lab Introduction
2
14th Jan.
Experiment #1 – Cation Analysis
3
21st Jan.
Experiment #2 – Anion Analysis
4
28th Jan.
Experiment #3 – Solubility Determination
5
4th Feb.
Tutorial – Ksp Problem Set
6
11th Feb.
Experiment #4 – pH Indicators and their Inflection Ranges
7
18th Feb.
FAMILY DAY
8
25th Feb.
Experiment #5 – Buffer Systems and Buffer Capacities
4th Mar.
STUDY BREAK
9
11th Mar.
Tutorial – Acid-Base Problem Set
10
18th Mar.
Experiment #6 – Inorganic Redox Reactions
11
25th Mar.
Experiment #7 – Microbiological pH Analysis
12
1st Apr
No LAB
13
8th Apr
Experiment #8 – Microbiological Redox Analysis
Experiment #7 –Follow up
14
15th Apr
Experiment #9 – Microbiological pH & Redox Analyses- TSI Test
Experiment #8 – Follow up
15
22nd Apr
Experiment #9 – Follow up
POLICIES
All laboratory reports are due at the beginning of the following lab period unless the instructor indicates otherwise. The lab instructor, at the beginning of each lab, will detail the write-up requirements. All laboratories must be completed before the end of week 14. Late reports will have a penalty of 20% per week.
Attendance in laboratories is mandatory. Lab reports will not be accepted for missed labs. For labs that missed due to illness, please provide a doctor’s note and make arrangements for the missed work with your instructor.
Experiment #1: Cation Analysis
Introduction:
Chemical analyses are categorized into two specific groups, namely qualitative analysis and quantitative analysis. The purpose of a qualitative analysis is to determine the identity or composition of a compound or mixture of compounds, whereas the goal in quantitative analysis is to determine how much of a given substance is present.
Qualitative analysis of an aqueous mixture typically proceeds in three stages:
1) The ions present in solution are separated into broad groups based on their solubility properties.
2) The individual ions within each group are separated by specific extraction techniques.
3) The identity of the specific ions are confirmed by specific chemical tests.
It is important to note that elements are grouped together in a qualitative analysis based on their solubility properties. Although these are called Group I elements, Group II elements, etc., this terminology should not be confused with the Group nomenclature within the periodic table, which is based on electronic configuration.
In cation analysis, typical qualitative analysis separates the ions present in a solution into six groups as shown in Table 1. Subsequent separation and identification is made possible by careful manipulation of the following:
i) the pH of the solution, ii) the known solubilities of different compounds of the ions at different conditions, iii) specific tests that produce a coloured compound with a particular ion, where the colour observed is characteristic of that ion only.
When conducting a complex qualitative analysis, the order of addition of reagents is important. Analysts commonly proceed with separating and testing for Group I ions, followed by Group II, Group III, and so forth. When conducting this type of analysis, one must ensure that all reactions have proceeded to completion, and that any preceding Group ions remaining in a solution are at such a low concentration that they will not interfere with subsequent tests. For example, a white precipitate formed in Group V cannot be AgCl, since all the Ag+ ions were precipitated in the Group I analysis.
Table 1: Qualitative Cation Analysis Chart
Group
Ions that can be precipitated
Precipitating reagent
Nature of precipitate
I
Ag+, Pb2+, Hg22+
hydrochloric acid
chlorides
II
(A) Hg2+, Cu2+, Cd2+, Bi3+, (Pb2+)
(B) As3+, Sb3+, Sn2+, Sn4+
hydrogen sulphide in presence of H+
hydrogen sulphide in presence of H+
sulphides - insoluble in ammonium sulphide
sulphides - soluble in ammonium sulphide
III
Fe3+, Al3+, Cr3+
ammonium hydroxide in presence of ammonium chloride
hydroxides
IV
Zn2+, Mn2+, Ni2+, Co2+
hydrogen sulphide in presence of OH-
sulphides
V
Ba2+, Sr2+, Ca2+, Mg2+
ammonium carbonate
carbonates
VI
Na+, K+, NH4+
specific tests required
------------
Experiment #1:
Separation & Identification of Group I Cations (Silver Group)
1. Add 5 drops of Group I stock solution to a glass test tube. 2. Add 5 drops distilled water to test tube. 3. Add 2 drops 6M HCl and mix. Insoluble chloride precipitates of silver, mercury (I) and lead will develop. 4. Wash sides of test tube with distilled water until precipitate is below level of solution. 5. Centrifuge test tube (be sure to balance test tube in centrifuge). 6. Test for complete precipitation by adding 1 drop of 6M HCl to the clear supernatant. If a precipitate forms, repeat step #5. 7. Place test tube in a cold water bath to chill sample. Then pour the clear solution above the precipitate out into another test tube. This process is known as decant and the clear solution collected is known as decantate. 8. Wash the precipitate from step #7 by adding 5 drops of distilled water and mix. Centrifuge sample, chill and decant. Put the decantate into the same test tube containing the decantate collected in step #7. The decantate may contain cations from the other groups and should be saved only if other analyses are to be conducted. In this case the decantate can be discarded. Save precipitate since it contains AgCl, PbCl2 and Hg2Cl2.
Separation of Lead from Silver and Mercury(I) & Confirmation of Lead:
To separate lead chloride from silver and mercury(I) chloride precipitates, analysts make use of the fact that PbCl2 is soluble in hot water while AgCl and Hg2Cl2 are not. After separation, the presence of lead ions are confirmed by adding potassium chromate to form the bright yellow insoluble lead chromate precipitate, or by adding sulphuric acid to form a white insoluble lead sulphate precipitate.
9. Add 10 to 15 drop of distilled water to the precipitate saved in step #8 and mix.
10. Place test tube in a hot water bath and allow the test tube solution to become hot. Stir solution frequently. (This is done to dissolve PbCl2.)
11. Centrifuge solution at once and decant as soon as possible. The decantate contains the soluble PbCl2, and the precipitate contains the insoluble AgCl and Hg2Cl2. Save both the decantate and the precipitate.
12. To confirm the presence of Pb2+, cool the decantate and divide into two equal portions.
To one portion, add 1 drop of 0.2M potassium chromate. If the bright yellow PbCrO4 precipitate forms, this confirms the presence of Pb2+.
To the second portion, add 1 drop of 2M sulphuric acid and wait (this reaction may be slow). If the white PbSO4 precipitate forms, this confirms the presence of Pb2+.
Separation and Detection of Silver and Mercury(I)
The separation of insoluble silver and mercury(I) chlorides is based on the fact that AgCl is soluble in ammonium hydroxide while Hg2Cl2 is not. Hg2Cl2 reacts with ammonium hydroxide to form the combined products of Hg and HgNH2Cl , both of which are insoluble. (Note: the presence of the greyish Hg and HgNH2Cl precipitate confirms the presence of Hg22+.)
The presence of Ag+ is confirmed by acidifying the solution of Ag+ in NH4OH and observing the precipitation of AgCl.
13. Add 4 drops of 15M NH4OH to the precipitate in step #11, mix well, centrifuge and decant. The decantate contains the soluble Ag+ ions. The presence of a grey or black-gray precipitate confirms the presence of Hg22+.
14. To the decantate in step #13, add 16M HNO3 dropwise, until the solution turns slightly acidic (confirm with litmus paper). A white precipitate confirms the presence of silver.
15. REPEAT STEPS 1 TO 15, SUBSTITUTING YOUR UNKNOWN FOR THE STOCK SOLUTION IN STEP #1. (Use only five drops of your unknown just like the stock solution. Your unknown may contain 1, 2 or 3 cations.))
Analysis Scheme for Group I Cations (Silver Group)
Ag+, Pb2+, Hg22+
(+ 6M HCl)
Decantate
ppt: AgCl PbCl2 Hg2Cl2 Cations of Groups II through VI colour: white white white
(+hot water)
Decantate
ppt: AgCl Hg2Cl2 Pb2+ colour: white white
(+H2SO4) (+K2CrO4)
ppt: PbSO4 PbCrO4 colour: white yellow Pb2+ present
(+NH4OH)
Decantate ppt: Hg + HgNH2Cl Ag(NH3)2Cl colour: grey
Hg22+ present
(+HNO3)
ppt: AgCl colour: white
Ag+ present
Analysis Scheme for Group I Cations (Known Solution)
Group I Cations
______, ______, ______
(+ 6M HCl)
Decantate
ppt: Yes or No ppt may contain: ______ ______ ______ Colour: ______ colour: ______ ______ ______
(+hot water)
Decantate ppt: Yes or No ppt may contain: ______ ______ Colour: ______ colour: ______ ______ May contain:______
(+H2SO4) (+K2CrO4)
ppt: ______ ______ colour: ______ ______
______ present
(+NH4OH)
Decantate ppt: ______ Colour:______ colour: ______ May contain: ______
______ present (+HNO3)
ppt: ______ colour: ______ ______ present
Analysis Scheme for Group I Cations (Unknown Solution #______)
Group I Cations
______, ______, ______
(+ 6M HCl)
Decantate
ppt: Yes or No ppt may contain: ______ ______ ______ Colour: ______ colour: ______ ______ ______
(+hot water)
Decantate ppt: Yes or No ppt may contain: ______ ______ Colour: ______ colour: ______ ______ May contain:______
(+H2SO4) (+K2CrO4)
ppt: ______ ______ colour: ______ ______
______ present
(+NH4OH)
Decantate ppt: ______ Colour:______ colour: ______ May contain: ______
______ present
(+HNO3)
ppt: ______ colour: ______
______ present
Experiment #2: Anion Analysis
Introduction:
As with cation analyses, anion analyses are also categorized into two specific groups previously discussed, namely qualitative analysis and quantitative analysis. However, each anion is tested for individually based on a sample taken from the original unknown solution.
For this experiment, specific anion tests for the qualitative determination of carbonates, chlorides, sulphates, nitrates and phosphates are provided. The specific chemistry of each test is as follows:
Carbonates: CO32- (aq) + 2H+ (aq) CO2 (g) + H2O
Chlorides: Cl- (aq) +Ag+ (aq) AgCl (s) (white ppt)
Sulphates: SO42- (aq) + BaCl2 (aq) BaSO4 (s) + 2Cl- (aq) (white ppt)
Nitrates: 8Fe2+ (aq) + 2NO3- (aq) + 8H+ (aq) 2Fe(NO)2+ + 6Fe3+ (aq) + 4H2O
(Note: Reduction of NO3- to NO by iron(II) results in formation of brown Fe(NO)2+ complex)
Phosphates: PO43- (aq) + 12MoO42- (aq) + 24H+ (aq) + 3 NH4+ (aq) (NH4)3 PO4.12MoO4 (s) + 12 H2O (yellow ppt)
Procedures:
Carbonates: Place about 20 drops volume of aqueous carbonate solution or solid mixture containing carbonates into a test tube. Add 3 to 4 drops of 6M HCl and observe. If a rapid effervescence occurs with a solid, or if bubbles form in an aqueous solution, this may indicate the formation of CO2 gas, indicating the presence of carbonates in the unknown. (If one wants to confirm the gas produced is carbon dioxide, allow the gas to come in contact with a saturated solution of Ba(OH)2 or Ca(OH)2. The formation of a white precipitate will confirm the gas produced is carbon dioxide.)
Chlorides: Place about 20 drops of solution in a test tube, then add 3 to 4 drops of 3M HNO3 until the solution is slightly acidic. Add 1 or 2 drops of 0.2M AgNO3 and mix. If a white precipitate forms, which is subsequently soluble in ammonium hydroxide, then Cl- ions were present in the original sample.
Sulphates: Place 20 drops of solution in a test tube, then add 3 to 4drops of 3M HNO3 until the solution is slightly acidic. Add 2 to 3 drops of 0.2 M BaCl2 and mix. If a white precipitate forms, then SO42- ions were present in the original sample.
Nitrates: Place 20 drops of sample into a test tube, then add a few crystals of solid ferrous sulphate or a few drops of aqueous ferrous sulphate solution. Carefully add about 2 mL of concentrated sulphuric acid by pouring it down the wall of the test tube while test tube is at a 45o angle. When this step has been completed, there will be two layers of solution in the test tube. The appearance of a brown ring at the interface between the two solutions indicates the presence of the NO3- ion in the original sample.
Phosphates: Place 20 drops of sample in a test tube, then add 4 to 5 drops of 3M nitric acid and 3 to 4 drops of 0.2 M ammonium molybdate solution. Place test tube in a hot water bath and observe. The formation of a slowly forming yellow-coloured solution or precipitate indicates the presence of the PO43- ion in the sample.
REPEAT ALL ANALYSES FOR THE UNKNOWN PROVIDED BY INSTRUCTOR.
Experiment #3: Solubility Determination
Introduction:
Many inorganic compounds are only slightly soluble in water. When these compounds are dissolved in water, the solution will become saturated with the cations and anions present in the compound.
i.e. CuCl(s) Cu+(aq) + Cl-(aq)
At this point, we can state that the ions in solution are in equilibrium with the solid, and the ion concentrations present in the water are at their greatest concentration. Addition of any more cations or anions will result in the precipitation of the slightly soluble solid.
The equilibrium solubility concentrations of both the cations and anion in solution may be expressed using a Solubility Product Constant. This constant may be determined for any slightly soluble compound having the form AxBy as follows:
For the solubility reaction: AxBy(s) x A+(aq) + y B-(aq)
Ksp = [A+(aq) ]x [B-(aq) ]y
where [ ] refers to the ion concentrations in units of molarity (mol/L).
The value of the Ksp constant is highly dependent on temperature, and is always reported at a specific temperature. (i.e. Ksp of CuCl = 3.2x10-7 @ 25oC). In addition, the solubility product constant indicated that it is the “product” of the cation and anion concentrations that determine if precipitation will occur, so the analyst must guard against looking at one ion independently of the other.
In this lab, the students will familiarize themselves with the concept of equilibrium solubility by determining the solubility of PbCl2 at various temperatures. From these data, the student will estimate the solubility product of PbCl2 and compare it to the literaure value.
Ksp of PbCl2 = [Pb2+][Cl-]2
[Pb2+] = molar solubility of PbCl2
[Cl-] = 2 x molar solubility of PbCl2
Procedure:
1. Prepare a boiling bath using 300 to 400 mL of tap water in a 600 mL beaker.
2. Using weighing paper, carefully weigh PbCl2 (to the nearest +/- 1 mg) in the amounts shown in the table below, and transfer into 5 marked, dry, large test tubes.
*Note: It is important that the test tubes are dry and you transfer the PbCl2 into the test tubes first before you add the water.
Tube # Sample weight Solubility Molar Solubility Temp. (g) (g/100g) (mol/L) (oC)
#1 0.500
#2 0.450
#3 0.400
#4 0.350
#5 0.300
3. Add 25.0 mL distilled water to each of the samples in the tubes, mix well, and mark the liquid level on the test tube. Place all test tubes in the boiling water bath ensuring that the liquid level in the test tube is below the liquid level of the boiling water bath.
4. When any solution has completely dissolved (this may require some mixing using a stirring rod), remove the numbered test tube from the boiling water bath, quickly check to make sure the liquid level in the test tube has not changed. If the level has dropped, bring the level back to the original level with distilled water. Insert a thermometer into the solution. and measure the temperature at which the first crystals begin to precipitate. Record the temperature.
5. Repeat step #4 for all solutions. (Note that Tube solution #5 will need to be cooled under cold running water.)
Analysis of Data:
Once all 5 crystallisation temperatures have been found, tabulate your results using the Table shown above . This includes the conversion of your solubility data into g PbCl2/100 g solution (assuming a solution density of 1.00 g/mL). Plot Molar Solubility of PbCl2 sol. vs. Temperature in oC .
Determine the molar solubility of PbCl2 @ 25oC from the graph. Calculate the Ksp (solubility product) at 25oC from the molar solubility and compare this value to the literature values which is 1.6 x 10-5.
Experiment #4: pH Indicators and their Inflection Ranges
Introduction:
In any biological or microbiological system, the pH of the solution is an important factor as to whether the species will be able to survive in its environment. (For example, when discharging water to Lake Ontario, the pH must be between 6.5 and 8.5, otherwise the fish population may be adversely affected.) The pH of any solution is defined by the mathematical equation pH = -log[H+]. To measure the pH of a solution, an analyst may use either a pH meter or a variety of pH indicators.
This lab introduces the concept of pH and indicators commonly used in chemistry and microbiology. Students will determine inflection ranges of known and unknown indicators, as well as using pH meters. They will compare their results with data reported in the literature.
Procedure:
Using the 4 common pH indicators, litmus, brilliant green, phenolphthalein and thymol blue, determine the inflection range and colour changes observed for each of the indicators. For consistency, when you record your results, always write down the colour change from acid to base and NOT in reverse. 1. Pre-test: Start with 2 test tubes, each 1/3 full with distilled water and add 1 drop of one indicator to each tube. Add 1 drop of 0.1 M HCl to one test tube and 1 drop of 0.1 M NaOH to the other tube and observe the colour changes, if any. If no colour changes occur, add 1drop of 1 M HCl and 1 M NaOH respectively to the same tubes and observe colour changes. Repeat for the other indicators.
2. Place 50 mL of distilled water in a 100 mL beaker. Using a stirring rod for mixing, add 8-10 drops of one indicator and record the solution colour. Depending upon your results in step #1, add 1 drop of either 0.1 M HCl or 0.1M NaOH and observe any colour change, if any. If no colour change is apparent, continue to add more drops of acid or base until the colour changes, but keep track of the number of drops added. (Note: If the pre-test indicated the use of a strong acid or base, use this instead.) Discard the solution and repeat for the other indicators.
3. Test: Once the approximate number of drop required is known, again place 50 mL of distilled water with 8-10 drops of indicator in a 100 mL beaker on top of a magnetic stir plate. Put a stir bar and insert a standardized pH probe into the beaker. The starting pH should be about 6. Turn stirrer on by using the left knob. (Do not turn on heat.) Add the appropriate number of drops of acid or base required for colour changes as determined in the previous step. Record the pH for the colour change. For best results, use fractions of drops. Always use the squeeze bottle or swirl the contents to wash down reagent drops adhering to the beaker walls. Obtain the precise upper and lower pH values at which the indicators no longer change colour. (If you overshot the colour change, you can go back by using acid or base.) Repeat for the other indicators.
Note: Before using the pH meter, standardize it. Set it to 25oC, then use a buffer solution with a constant pH, e.g. pH = 7.00. Rinse the electrode after switching to Stand-by. Always store the electrode in buffer solution. To check the pH of a solution, immerse the electrode into the solution after rinsing the probe with distilled water. Swirl solution for approximately 20 to 30 seconds and then take the pH reading.
Note: Thymol Blue has 2 inflection ranges.
4. Record colour changes and pH range of the inflection point.
5. Determine the identity of an unknown indicator by repeating steps #1 to 3 as previously discussed. Base the identity on both colour changes and pH inflection point data. Record
LIST OF COMMONLY USED pH INDICATORS
Name pH Range Colour (acidic to basic)
Crystal Violet 0.0 - 1.8 yellow to blue
* Brilliant Green yellow to green
Malachite Green 1.0 - 1.8 yellow green to blue green
*Thymol Blue red to yellow
Crystal Violet 1.6 - 2.5 green to violet
Methyl Violet 1.8 - 2.5 violet to blue
Methyl Orange 3.2 - 4.4 pink to yellow
Bromophenol Blue 3.0 - 4.6 slight yellow to blue
Congo Red 3.0 - 5.0 violet to red/orange
Bromocresol Green 3.8 - 5.4 yellow to blue
Methyl Red 4.8 - 6.0 pink to yellow
Mod. Methyl Red 4.3 - 5.8 violet to blue/green
* Litmus red to blue
Bromocresol Purple 5.2 - 6.8 slight yellow to purple
Bromothymol Blue 6.0 - 7.6 yellow/orange-green-blue
Neutral Red 6.8 - 8.0 purple to slight yellow
Phenol Red 6.6 - 8.0 violet to orange/red
Thymol Blue yellow to blue m-Nitrophenol 6.8 - 8.6 colourless to yellow
* Phenolphthalein colourless to pink
Thymolphthalein 9.4 - 10.6 colourless to blue
Alizarin Yellow 10.0-12.0 clear to deep yellow
Known indicators used in the experiment. pH inflection range not reported.
Experiment #5: Buffer Systems and Buffer Capacities
Introduction:
A buffer solution is an equilibrium system which, within limits, resists changes to pH upon the addition of strong acids or strong bases. Buffer systems are made when a weak acid or weak base is mixed with its corresponding salt. Buffers are commonly incorporated into microbiological growth media to ensure the pH range of the solution remains within the acceptable pH range for biological growth.
The capacity of a buffer is the amount of acid or base required to change the pH of 1 litre of buffer by one pH unit. It is expressed in units of moles of acid or base per litre buffer. Depending upon the pH, acidic buffers are titrated with a base and basic buffers are titrated with acid.
This lab familiarizes the student with the concept of commonly used buffer systems. Students will produce buffers for specified pH and perform buffering capacity measurements and compare their results with literature data.
Procedure:
1. Prepare a Butterworth Buffer from the table below according to the pH specified by the instructor. Using a volumetric cylinder and a Mohr pipette, the appropriate volumes of the two buffer components are mixed and measured using the previously calibrated pH meter. The measured pH should agree within +/- 0.2 pH units. Record results.
Butterworth Buffer Table for pH 5.80 to pH 8.00 Buffers pH x pH x
5.80 3.6 7.00 29.1
5.90 4.6 7.10 32.1
6.00 5.6 7.20 34.7
6.10 6.8 7.30 37.0
6.20 8.1 7.40 39.1
6.30 9.7 7.50 41.1
6.40 11.6 7.60 42.8
6.50 13.9 7.70 44.2
6.60 16.4 7.80 45.3
6.70 19.3 7.90 46.1
6.80 22.4 8.00 46.7
6.90 25.9
Note: Buffer solution made using 50 mL of 0.1M KH2PO4 and x mL of 0.1M NaOH. Final volume of all buffer mixtures is 100 mL.
Since the given NaOH solution is not exactly 0.1 M it is necessary to calculate and use the equivalent volume.
2. Once the buffer has been satisfactorily prepared and shown to the instructor, the student proceeds with the buffer capacity measurement. Using a pH meter, stirrer bar and magnetic stirrer plate, and a burette mounted on a stand, fill the burette to exactly zero mL with the given standardized 0.1M HCl. The lower meniscus is to be taken as your liquid level. Use 50 mL of your buffer solution, as measured with a graduated cylinder. The remaining 50 mL of your buffer solution is use for step #4.
3. Add reagent slowly, drop by drop,(one drop is approximately 0.05 mL) from your burette until a pH change of 0.1 units has been observed. Read your reagent consumption from your burette and record it. Continue process in intervals of 0.1 pH units, until the buffer system’s pH has changed by 1.1 units.
4. Repeat steps #2 and 3 with the remaining 50 mL of your buffer solution, but change the reagent in the burette to the given standardized 0.1M NaOH.
5. From the volume of HCl and NaOH added find the # of mol of H+ and OH- added for each data point. Plot pH vs # of mol of H+ and # of mol of OH- as shown in the figure below.
6. From the plot obtain the # of mol of H+ and OH- required to change the pH by 1.0 pH unit for 50 mL of your buffer solution. Calculate your nominal buffer capacity for a change in 1.0 pH units in # of mol of H+ (acid) and OH- (base) per litre of buffer solution.
pH
# mol of H+ added 0 # of mol of OH- added
Plot of pH vs mol of HCl and mol of NaOH.
Experiment #6: Inorganic Redox Reactions
Introduction:
Biochemical redox (reduction-oxidation) reactions are commonly used in microbiology to identify microorganisms. These reactions are often based on metabolic products formed by specific microorganisms, which can then be subsequently identified. For example, nitrate reducing bacteria may convert the nitrate ions (NO3-) present in water to nitrite ions (NO2-) under anaerobic conditions. To identify the nitrate reducing bacteria, an analyst would inoculate a nitrate broth with bacteria and subsequently monitor the solution for the formation of nitrites.
At other times, specific compounds are added to a growth medium which can only be utilized by certain micro-organisms which are often growth inhibiting (toxic) to other microorganisms. One common example of these types of compounds is oxygen which separates aerobes from the anaerobes.
The purpose of this laboratory is to familiarize the student with the concept of redox reactions, and introduce the student to typical redox indicators used in microbiological analyses.
Procedure:
1. Methylene Blue Redox Test
A. Fill a test tube 1/4 full with distilled water. Add three to five drops of methylene blue, three drops of 6 M hydrochloric acid and a pinch of Zn dust. The evolution of hydrogen gas reduces the methylene blue dye to a colourless reduction product.
Chemical Reaction: 2HCl(aq) + Zn(s) H2(g) + ZnCl2(aq)
Redox Reactions: Zn Zn2+ + 2e-
2H+ + 2e- H2 methylene blue (oxidized) + 2e- methylene blue (reduced)
(blue) (colourless)
B. Decant the clear solution from the zinc dust by pouring it slowly in another test tube. Add 2 drops of 3% hydrogen peroxide (H2O2) solution. Record your results.
Reaction: methylene blue (reduced) + H2O2 + 2H+ methylene blue (oxidized)+
2H2O
2. Nitrate Reduction Test:
Using 4 test tubes, each filled to ¼ full with distilled water, do the following tests and record your observations.
Tube #1: 3 drops Nitrate solution(1%) + 3 drops Nitrite ID agent + 1 drop 5M HCl
KNO3 + nitrite indicator + HCl no reaction
Tube #2: 3 drops Nitrate solution(1%) + 3 drops Nitrite ID agent + 1 drop 5M HCl + Zn
KNO3 + 2HCl + Zn ZnCl2 + KNO2 + H2O
KNO2 + nitrite indicator azodye (pink/red)
Tube #3: 3 drops Nitrite solution(1%) +3 drops Nitrite ID agent + 1 drop 5M HCl
KNO2 + nitrite indicator azodye (pink/red)
Tube #4: 3 drops Nitrite solution(1%) +3 drops Nitrite ID agent + 3 drop 5M HCl + Zn
NaNO2 + 3Zn + 8HCl NH4Cl + 3ZnCl2 + 2H2O + NaCl
The pink-red colour observed in tubes #2 and 3 is based on the formation of an azodye, which is formed when nitrite is present.
Note: Heating the test tubes may be necessary to speed up reaction. If required, place test tubes in a hot water bath.
3. Hydrogen sulfide formation from thiosulphate:
Into each of 3 test tubes, each filled to ¼ full with distilled water, add a few drops of sodium thiosulphate solution(5%) + ferrous sulfate solution(5%). Add the following as shown, and record your observations.
Tube #1: Add a pinch of zinc dust plus 1 drop dilute HCl.
Na2S2O3 + 3Zn + 8HCl H2S + 3ZnCl2 + S0 + 3H2O + 2 NaCl
Redox Reactions: S2O32- + 6H+ + 6e- S0 + S2- + 3H2O
Zn Zn2+ + 2e-
Test for evolution of H2S by gently smelling the contents of the test tube. Hydrogen sulphide has a characteristic “rotten egg” smell.
Proof for H2S formation: FeSO4 + H2S + 2NH3 FeS (black ppt) + (NH4)2SO4 Tube #2: Add only a pinch of zinc dust:
Na2S2O3 + Zn no reaction
Tube #3: Add only 1 drop dilute HCl:
Na2S2O3 + HCl S0 + H2SO3 + 2NaCl
S2O32- S0 (yellow/white ppt) + SO32-
4. Testing Redox Agents with Starch and Iodine/Iodide: Fill 2 test tubes to ¼ full with room temperature distilled water and 1 test tube with cold distilled water. Add 2 drops of 10% acetic acid to each of the test tubes. Do the following tests and record your results.
Tube #1: Add 3 drops of iodine (KI-I2) solution plus 3 drops of starch solution to one of the test tubes with room temperature water. Note the purple to blue colour formed. The starch is being used to indicate the presence of iodine.
Tube #2: Add 3 drop of potassium iodide (KI) solution plus 3 drops of starch solution to the test tube with cold distilled water. Record initial observations. Shake the tube vigorously for 1-2 minutes and observe again. You should see a slight blue colouration after a while, based on atmospheric oxidation of iodide to elemental iodine. (This procedure works better with cold water.)
2H+ + ½O2 + 2I- I2 + H2O
Redox Reactions: 2H+ + ½O2 + 2e- H2O
2I- I2 + 2e-
Tube #3: Add 3 drops of iodine (KI-I2) solution plus 3 drops of starch to the test tube with room temperature distilled water. Observe, then add several drops of thiosulphate (Na2S2O3) solution and observe again.
2Na2S2O3 + KI-I2 + 2HAc Na2S4O6 + KI + 2HI + 2NaAc
As in the previous cases, starch is used as an indicator for the presence of iodine. When sodium thiosulphate is added, the iodine is reduced to iodide, and will not test positive with the starch test. Note: This reaction is used in quantitative analytical iodometric titrations (i.e sulphite tests when determining presence of sulphate and sulphite reducing bacteria).
Experiment #7: Microbiological pH Analysis
Introduction:
In many microbiological systems, bacteria will ferment a carbohydrate broth to produce various acidic byproducts. Depending upon the carbohydrate used, the initial pH of the broth, presence of various organic and/or inorganic compounds, the type of bacteria, etc., the identification of the microorganism may be determined. For example, in the mannitol salt medium, only bacteria that ferment the sugar mannitol in a concentrated salt broth will survive (typically staphylococci). By using a pH indicator to detect the presence of acidic byproducts, a microbiologist may identify the presence of staphylococci using this media.
In this experiment, the student will prepare various phenol red carbohydrate media and apply his/her knowledge of pH indicators (Expt. #4) in determining the function of various simple microbiological media.
Procedure:
Each group of students will be assigned a broth to prepare and share with the other groups. Each broth will have a volume of 250 mL and will consist of the following:
Ingredients
Base Broth
Mannitol Broth
Lactose Broth
Maltose Broth
Peptone
2.5g
2.5g
2.5g
2.5g
Beef Extract
0.25g
0.25g
0.25g
0.25g
Mannitol
1.25g
Lactose
1.25g
Maltose
1.25g
Sodium Chloride
1.25g
1.25g
1.25g
1.25g
Phenol Red
2 mL
2 mL
2 mL
2 mL
Water
to 250 mL
to 250 mL
to 250 mL
to 250 mL Gently boil the medium for 10 minutes to sterilize the medium and drive off any dissolved gases. Wait until the medium has cooled down sufficiently adjust the pH of each broth to 7.4 ± 0.2 by adding 0.1 M.NaOH or 0.1 M HCl dropwise. Pour medium into culture vials, and add a Durham tube to each vial to assess if gas formation occurs. Label each vial to indicate the type of broth and give one vial to each of the other groups.
Record initial condition and colour of vial. Record final condition(s) and colour during next laboratory session.
NOTE: Mark your individual group cultures carefully. Each culture will be inoculated with bacteria after being autoclaved. They will be incubated in the microbiology laboratory. Incubated vials will be available to students at the next scheduled laboratory session for observations.
Post-Laboratory Assignment:
1. Based on your knowledge of microorganisms, briefly discuss whether the bacteria used to inoculate each broth were able to ferment the carbohydrates present in the broth. On what basis are these statements made? 2. What is the purpose of each ingredient in the broth mixture assigned to your group?
3. If phenol red was not available for preparing the medium, what suitable pH indicator(s) may be used? How would this impact on the test?
Experiment #8: Microbiological Redox Analysis
Introduction:
Some bacteria, such as Salmonella sp. and Proteus sp., are able to produce hydrogen sulphide (H2S) by reducing inorganic compounds such as thiosulphate (S2O32-), sulphates, (SO42-) , and/or sulphites (SO32-). In each case, the oxidation state of sulphur changes from the +2, +6 and +4 respectively to the -2 oxidation state in the acidic conditions associated with the bacteria. To use this fact as a basis of determination, the SIM medium is used to promote growth of bacterial strains that are capable of reducing thiosulphate to hydrogen sulphide.
In this experiment, the student will prepare a SIM medium, and apply his/her knowledge of redox indicators (Expt. #6) in determining the function of this microbiological media.
Procedure:
Each group of students will prepare 200 mL of SIM medium as follows:
Ingredients
SIM Media
Peptone
6.0g
Beef Extract
0.6g
Ferrous Sulphate
0.04g
Sodium Thiosulphate
0.005g
Agar
0.7g
Water
to 200 mL Gently boil the medium for 10 minutes to sterilize the medium and drive off any dissolved gases. Wait until the medium has cooled down sufficiently adjust the pH of each broth to 7.3 ± 0.2 by adding 0.1 M.NaOH or 0.1 M HCl dropwise. Pour media into four (4) culture vials to about half full and allow the medium to solidify in the vertical position. Label each vial to indicate the name of the group and sample number.
Record initial condition and colour of vial. Record final condition(s) and colour during next laboratory session.
NOTE: Mark your individual group cultures carefully. Each culture will be inoculated (stab) with bacteria after being autoclaved. They will be incubated in the microbiology laboratory. Incubated vials will be available to students at the next scheduled laboratory session for observations.
Post-Laboratory Assignment:
1. Based on your knowledge of microorganisms, briefly discuss whether the bacteria used to inoculate each broth were able to reduce the thiosulphate present in the medium? On what basis are these statements made? 2. What is the purpose of each ingredient in the medium?
Experiment #9: Microbiological pH & Redox Analyses - TSI Test
Introduction:
When provided with a sample of unknown bacteria, an analyst will often rely on test media that provide information on both the types of carbohydrates a bacteria can ferment, as well as the microorganism’s ability to reduce various inorganic compounds, as an aid in the identification process. The Triple Sugar-Iron Agar test (TSI) is a rapid screening test that is able to differentiate and distinguish the members of the Enterobacteriaceae and other intestinal bacilli (intestinal bacteria found in contaminated food and water supplies).
The TSI medium utilizes the three sugars at the following concentration - lactose (1%), sucrose (1%) and glucose (0.1%), as well as sodium thiosulphate to provide differentiation among the various bacteria.
In this experiment, the student will prepare a TSI medium and apply his/her knowledge of pH and redox indicators in determining the function of this microbiological media.
Procedure:
Each group of students will prepare 200 mL of TSI media as follows:
Ingredients
TSI Media
Ingredients
TSI Media
Beef Extract
0.6g
Sucrose
2.0g
Yeast Extract
0.6g
Ferrous Sulphate
0.04g
Peptone
3.0g
Sodium Chloride
1.0g
Proteose Peptone
1.0g
Sodium Thiosulphate
0.06g
Dextrose
0.2g
Agar
2.4g
Lactose
2.0g
Phenol Red
5 mL Gently boil the medium for 10 minutes to sterilize the medium and drive off any dissolved gases. Wait until the medium has cooled down sufficiently adjust the pH of each broth to 7.4 ± 0.2 by adding 0.1 M.NaOH or 0.1 M HCl dropwise. Pour medium into four (4) culture vials to about half full. Label each vial to indicate the name of the group and sample number.
Record initial condition and colour of vial. Record final condition(s) and colour during next laboratory session.
NOTE: Mark your individual group cultures carefully. Each culture will be inoculated (stab and streak) with bacteria after being autoclaved and allowed to solidified in a slant position. They will be incubated in the microbiology laboratory. Incubated vials will be available to students at the next scheduled laboratory session for observations.
Post-Laboratory Assignment:
1. Based on your knowledge of microorganisms, briefly discuss whether the bacteria used to inoculate each broth were able to ferment the carbohydrates and reduce the thiosulphate present in the medium? On what basis are these statements made? 2. What is the purpose of each ingredient in the medium?
3. Originally, this test utilized bromothymol blue (inflection point: 6.0-7.6) as the pH indicator. How would this impact on the test?
4. In the original industrial production of penicillin, the strain Penicillium chrysogenum was introduced into a culture medium consisting of the following:
Grams per Litre Medium
Corn liquor 30.0 g
Lactose 30.0 g
Glucose 5.0 g
NaNO3 3.0 g
MgSO4 0.25 g
ZnSO4 0.044 g
Phenyl acetamide 0.05 g
(penicillin precursor)
CaCO3 3.0 g
NaHCO3 3.0 g
The bacterial strain would utilize the food and nutrients in the media, in the presence of oxygen from air bubbled through the mixture, to produce penicillin. The penicillin produced was subsequently extracted and purified for use.
a) What is the purpose of each of the ingredients present in the medium?
b) Calculate the pH of the medium based on the presence of the carbonate ion in equilibrium with the bicarbonate ion according to the reaction:
HCO3- H+ + CO32- Ka2 = 5.61 x 10-11
Hint: Assume all the sodium bicarbonate in the medium is soluble. Determine the concentration of the carbonate ions based on the solubility product of CaCO3 .
c) If the laboratory did not have ZnSO4 available, could it be replaced with ZnS on an equivalent molar basis? What concerns would you have?
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In a lab, you must always make safety a priority in case of an emergency. In our lab, we made sure to wear closed-toed shoes as well as goggles at all times. We also performed part of our experiment in a fume hood in order to protect ourselves from any potentially dangerous fumes. We made sure not to touch the hotplate, and also to use tongs when picking up a hot beaker.…
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If safety equipment is provided, then staff must ensure it is used at all times e.g if gloves are provided for using hazardous materials in science, then they must be worn.…
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released by the breakdown of glucose is stored in the high energy phosphate bonds of…
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* Under no circumstances should food or drink be consumed in the lab area. You may drink in the classroom, but drinks are not to leave the desk section of the class.…
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Other safety concerns associated with this exercise are basic lab precautions, including keeping a clean workspace, washing your hand after dealing with chemicals and blood, and using precaution when obtaining blood samples. In order to ensure that this occurs, layout the equipment beforehand in a clean sterile environment and clean the work space after experiment. By doing so, this reduces the risk of contamination or cutting yourself when using the tools. Use gloves, masks, and goggles at all times.…
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Cited: Please note that in the introductory laboratory course, you will not be required to properly document sources of all of your information. One reason is that your major source of information is this website, and websites are inappropriate as primary sources. Second, it is problematic to provide a hundred students with equal access to potential reference materials. You may nevertheless find outside sources, and you should cite any articles that the instructor provides or that you find for yourself.…
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