chemistry practical book
DESCRIPTION
HSC Chemistry. List of practicalsTRANSCRIPT
Reactivity Of Alkanes And Alkenes Aim: To compare the reactivity of Cyclohexane and Cyclohexene using bromine water
Hypothesis: The cyclohexene will react due to the presence of its double bond.
SAFETY:
1. Wear safety goggles and gloves at all times during the experiment, as the organic solvent
can be absorbed by soft tissues.
2. Organic Liquids are volatile and flammable, Hence keep away from fire sources, Use very
small quantities of chemicals [Can cause skin irritations and respiratory problems]
3. Toxic Fumes, Avoid breathing in bromine water, performed in a fume cupboard.
4. WASTE DISPOSAL: Dispose of chemical wastes in supplied waste containers, not sink.
Method
1. Place two test tubes in a test tube rack and add 5 drops of bromine water to each teats
tube, observe and record their initial colour in a table.
2. Add 8 drops of Cyclohexene to the first test tube then stopper it, start the stopwatch and
then gently shake the test tube and place it back on the rack.
3. Stop the stopwatch at 2 min and observe and record the final colour of solution in a table
4. Repeat steps 2 and 3 using Cyclohexane in the second test tube
5. Dispose of waste chemicals in the waste containers supplied by the teacher.
6. Repeat the entire experiment 10 times.
Justify the method:
- Cyclohexene was used instead of ethylene or propene, as C1 to C4 are gases at room temp
- Cyclohexene and Cyclohexane were used instead of hexane/ane as cyclic hydrocarbons are
more stable than their linear counterparts, Less toxic etc
Results
Hydrocarbon Initial Colour Of Bromine Water Final Colour of solution
Cyclohexene Reddish-brown Colourless
Cyclohexane Reddish-brown Reddish-brown
DISCUSSION:
- The experiment showed that bromine water changed colours from reddish-brown to
colourless when cyclohexene was added to it. [Since alkenes have a highly reactive double
bond (High electron density), Which easily reacted with the bromine water [Br2(aq)].
- However when cyclohexane was added to bromine water no decolourisation occurred
o Since alkanes more stable single bonds, Need addition energy to break the C-H bonds, If
UV light present, reaction will undergo slow substitution reactions.
ACCURACY: A control was used (the test tubes with no bromine water)
- Accuracy affected by; The age of bromine water (Made colour changes less visible) Need
fresh Br2(aq), Dropper was aimed at centre of test tube, Same person used stopwatch for test
o Accuracy was affected by small quantities of soln| Same size of test-tubes were used
VALIDITY:
- Fairly valid, Only one variable was changed, All other Variables were controlled, Include ↓
o CONTROLLED: The number of drops of Cyclohexane/ene and bromine water used, All
reactions were stoppered after the reactants and was gently shaken for the same time.
o INDEPENDENT: The type of hydrocarbon whether it was an alkene or alkane.
o DEPENDENT: The rate of decolorisation of the bromine water from brown to clear.
- Validity was affected by the Amount of UV light in the room (Provided addition energy for
other reaction) Not near window, Minimised by covering in aluminium foil.
- Overall the experiment is valid as all necessary variables were controlled.
RELIABILITY:
- The experimental results are reliable because it was repeated several times and consistent
results were achieved, It also accounted for theoretically expected results, Hence reliable
IMPROVEMENTS:
- Using larger quantities of the solution, Freshly made bromine water
- Using a measuring cylinder to measure volumes of reactants
CONCLUSION:
Cyclohexene decolourises in bromine water hence is more reactive than cyclohexane which
does not decolourises in bromine water| Indicates alkenes more reactive than matching alkane
Fermentation Practical AIM: To carry out the fermentation of glucose and monitor mass changes.
EQUIPMENT:
o Electronic balance o 250ml Conical flask o Rubber stopper with hose
o 150ml beaker
o 20g of glucose (Sugar) o 7g of dried yeast o l00ml of water o 75ml limewater [Ca(OH)aq]
SAFETY: Wear safety glasses and gloves to avoid the spashing of chemical, I.e Prevent limewater from coming into contact with eyes since limewater is very basic (Ca(OH)2).
METHOD:
1. Measure 250 ml of water, 25 g of glucose, 7 g of yeast respectively using an electronic
balance, and add them into the conical flask and swirl to mix.
2. Measure and record the initial mass of the conical flask and its contents
3. Half fill a 150ml beaker with limewater and measure and record its mass using balance
4. Stopper the conical flask and completely immerse the end of the tube into the limewater
5. Put the two containers in a warm area and record the appearance mixture and limewater
6. Repeat the weighing and observations of each container daily for 2 more days.
7. Repeat the entire experiment 5 times.
Results
* Include a table column with change in mass
Discussion
- The experiment showed that the mass of the fermentation mixture decreased from x to
y, as CO2 Gas produced by fermentation
dissolved in the lime water, producing a
milky white precipitate, with the corresponding increase in the mass of limewater
ACCURACY:
- Fairly accurate, Electronic balance measured to ± 0.1g accuracy
- Environment was subject to changes in temperature
- Leaks were minimised by taping the conical flask to the rubber stopper.
VALIDITY:
- The experiment is valid as it investigates the aim. Also Only one variable was changed
(Independent), While the other Variables were controlled
o Independent: Number of days | Dependent: Mass Change
o Control: Location in room, Temperature of environment
- Validity affected as some of CO2 escaped from mixture, Different then theorectical.
RELIABILITY:
- The experimental results are reliable because were compared with the class and
consistent results were achieved
Conclusion
- The fermentation of glucose decreased the mass of the mixture by 5.9g and the
limewater turned milky, indicating the presence of C02.
Heat Of Combustion Practical AIM: To determine and compare the heats of combustion of ethanol, 1-butanol and l-
propanol per gram and per mole.
EQUIPMENT:
ethanol spirit burner
l-butanol spirit burner
I -propanol spirit burner
200m1 water
250m1 beaker
Stirring rod
Thermometer
Tripod
Gauze mat
Heat proof mat
SAFETY:
- Be careful with these highly flammable chemical and wipe any spills before lighting a match.
- Wear your safety goggles to prevent chemicals or hot water touching your eyes.
- Tie back long hair so it cannot catch fire
- Do not inhale poisonous l-propanol or l-butanol vapours.
Method
1. Pour 200ml of water into a beaker, then measure its mass using an electronic balance
2. Measure the mass of each spirit burner using an electronic balance and record mass
3. Set up apparatus as shown in the diagram. 4. Measure the temperature of the water using a thermometer and record it in a table 5. Light the ethanol spirit burner and stir the water while it is being heated. 6. Measure and record the temperature of the water after it has increased by l0oC. 7. Put out the spirit burner flame and let it cool before measuring its mass and recording it
in a table. Repeat steps 4 to 7 for each l-propanol and 1-butanol. 8. Calculate the heats of combustion for each liquid alkanol per gram and per mole.
Mass 200ml water = 189.6 |Also include amount of fuel burn (g)| Need molar mass for calculations
Calculations
DISCUSSION:
- The experiment showed that Butanol had the highest heat of combustion with X kJ/Mol,
Followed By Propanol with Y kJ/Mol, and finally ethanol with z kJ/Mol
o Larger Alkanols have extra CH2 group to chain, and thus extra H2O and CO2 are
produced, more bonds formed, > Energy released. Higher heat of combustion
- Accuracy: Very inaccurate, showed a X % error for the combustion with the theoretical
value. The calculated figure differ from the actual figure because:
o Heat lost to atmosphere, absorbed by tripod, gauze mat and beaker|Use Wind shield
o The thermometer touching the beaker instead of measuring the temp of the water.
o There may have been some incomplete combustion of the ethanol.
o The soot (carbon) that deposited on the beaker may have absorbed some of the heat.
The accuracy of the experimental method or apparatus could be improved by:
- Using a temperature probe and data logger
- Using a balance that measured to more decimal places
Validity
- validity is compromised due to large amount of heat lost to the surroundings
- The experiment method is partly valid as it is able to compare the Heat of combustion
between Alkanols, However did not accurately calculate heat of combustion
o Variable controlled- Independent: Type of Alkanols | Dependent: Heat of Combustion
The validity of the experimental method or apparatus could be improved by:
o Attaching the thermometer in the centre of the beaker, also removing the tripod and
gauze mat and holding the beaker up with a retort stand and clamp
o Moving the ethanol burner closer to the beaker of water, Using an insulated calorimeter
o Ensuring complete combustion of the ethanol by supplying sufficient oxygen
o Removing soot (carbon) from the bottom of beaker, Increase purity of Alkanols used
CONCLUSION: The heat of combustion of the Alkanols increased with molecular weight, It
was found that butanol had highest heat of combustion followed by Propanol and ethanol
Galvanic Cell Practical AIM: To investigate and measure the differences in voltage when different combinations of
metals are used in constructing a galvanic cell.
EQUIPEMENT:
o 3x l00mlbeakers
o 1 x2 ml beaker
o 1 x voltmeter
o 1 mol/L zinc nitrate, copper nitrate,
magnesium sulfate
o 3 x strips of filter paper (l cm by l0 cm)
o 3 x strips of filter paper (l cm by l0 cm)
o Strips of zinc, copper and magnesium
o Saturated potassium nitrate solution
o 2x connecting wires
o 2x alligator clips
o Steel wool
SAFETY:
- Lead nitrate is toxic, so hands must be washed after use, spills cleaned up immediately
- Wear safety goggles to prevent chemicals from coming into contact with eyes
- Dispose of waste solutions in a supplied waste container, do not pour down the sink.
METHOD: 1. Clean all the metal strips with steel wool.
2. Place 50ml of zinc nitrate, copper nitrate and magnesium sulfate soln into 3 separate beakers
3. Soak filter papers in saturated potassium nitrate solution and leave to soak.
4. Add the zinc to the zinc nitrate solution and the copper to the copper nitrate solution.
5. Take out a piece of filter paper and connect the zinc beaker to the copper beaker with the filter
paper dipping into each solution.
6. Connect the zinc and copper to a voltmeter using the wires and alligator clips.
7. Note which electrode is positive and record the magnitude of the voltage
8. Remove the salt bridge and note what happens.
9. Repeat steps 4 to 7 using the following metal combinations
Cell Beaker A Beaker B
1 Zn + 1M Zn (NO3) 2 Cu + 1M Cu (NO3) 2
2 Zn + 1M Zn (NO3) 2 Pb + 1M Pb (NO3) 2
3 Zn + 1M Zn (NO3) 2 Fe + 1M Fe (NO3) 2
4 Pb + 1M Pb (NO3) 2 Fe + 1M Fe (NO3) 2
5 Cu + 1M Cu (NO3) 2 Fe + 1M Fe (NO3) 2
6 Pb + 1M Pb (NO3) 2 Cu + 1M Cu (NO3) 2
Results
Discussion
- The experiment showed the voltage produced by metal combination, Also showed the
need of the salt bridge to complete the circuit (Balance the charges in each half-cell).
Reliability:
The experimental results are reliable because it was repeated several times and consistent
results were achieved, It also accounted for theoretically expected results, Hence reliable
Validity
- Valid, Only one variable was changed, All other Variables were controlled, Include ↓
o Controlled: Volume of solution used, type of salt bridge used and the voltmeter used.
o Independent: Metal Combinations | Dependent: Output voltage
Accuracy
- Not accurate as the value we obtained were significantly lower than theoretical results
o Since electrodes, alligator clips and wires were corroded ( Leads to > resistance,
reduced voltage), Analog Voltmeter Difficult to make measurements ± 0.1 V error
o Presence of impurities in solution, low quality voltmeters (High resistance)
o The salt bridge may have dried up, hindering the movement of ions reduce voltage
JUSTIFICATION
- We cleaned the electrode with steel wool before use to remove impurities
- We used a potassium nitrate salt bridge, which does not form precipitates with any of
the electrolytes used so does not affect the half-cell reactions help balance charges
Conclusion
- Each metal combination produced a voltage in their respective electrolyte solution
MODELING RECATIONS (using ball and stick models)
ADVANTAGES
- Allows us to represent bond breakage and bond formations that occur during reactions
- Provides a 3D representation of a reaction, which leads to a better understanding when
compared to liner equations |Shows the shapes of reactants and products
- Shows the spatial arrangement of atoms
DISADVANTAGES
- Ball and stick models cannot perfectly re-create the shape of reactants and products as
there are only certain places for sticks to attach to balls
- Hard to explain the strength of the bond.
THE ACIDIC ENVIROMENT- Indicator Prac
Aim: To prepare and test a natural indicator
Equipment:
Beetroot
Cutting Board +Knife
250mL beaker
l00ml water + Bunsen Burner
Dropper Bottle
5 x test tube
Test tube rack
Distilled Water
2mL of 1M Of NaOH
2mL of 2M Of NaOH
2ml of 1M HCI
2mL of 2M HCl
RISK ASSESMENT
- If acid/basic solution come into contact with skin and eyes, they may cause irritation and
burns, to avoid safety glasses and gloves must be worn at all times.
- During heating, the beaker may become hot, allow it to cool before handling it.
Method
1. Use a cutting board and knife to finely cut beetroot and place it in a beaker
2. Add 100 ml of water, and boil using the Bunsen burner for 5 min, and leave to cool
o Strain the beetroot solution to remove remaining beetroot
3. Add 5 drops of distilled water, 1M NaOH, 2M NaOH, 1M HCI and 2M HCI into 5 separate
clean test-tubes using dropper bottles
4. Using a dropper bottle draw up some of the beetroot solution (the natural indicator) and
add 5 drops to each test tube and record the change in colour of the indicator.
Results
In distilled water, it was DARK PURPLE.
In the NaCl solution, it was DARK RED.
In the HCl solution, it was PINK.
In the NaOH solution, it turned YELLOW
Discussion
- The experiment showed that the original purple colour changed to pink on addition of
HCL (Acidic Conditions) and yellow in NaOH (Basic Conditions)
Reliability: The experiment was reliable because it was repeated in exactly the same way by other groups in the class and we achieved similar results.
Accuracy: Results were qualitative and a beaker was used for measurement of the water
which was accurate enough to determine colour change. However a pH meter could be used
Validity
- The experiment is valid, Only one variable was changed, Other Variables were controlled,
o Control: volume of solution used, conc of solutions used, amount of indicator added.
o Independent: Acidic/basic nature of chemical |Dependent: Colour change indicator
JUSTIFICATION
- We also tested indicator with substances with a variety of pH to accurately determine
the colours, allowing us to test the colour of the indicator over large range of pH values.
- Beetroot was used as it readily available and gives off easily identifiable pigment
- Canned beetroot was not used as it often contains preservatives
LIMITATIONS: We were not able to determine the exact pH at which beetroot changes color
and the transition range of the indicator; this could have been done using a pH probe.
Conclusion:
Beetroot is a natural indicator that changes colour when mixed with an acid or base.
pH Probe Experiment Aim: To use pH probes and indicators to distinguish acidic, basic and neutral chemicals.
Aim 2: And To measure the pH of identical concentrations of strong and weak acids
Safety
- Wear safety glasses throughout this experiment.
- Dilute ammonia solution is mildly toxic. Its fumes should not be inhaled.
- Do not allow bleach to contact the skin, eyes or clothes.
- Identify other safety precautions relevant to this experiment by reading the method
Equipment and materials:
- Test Tube
- pH Meter + Buffer Solution
- Test Tube Rack
- Distilled Water
- 1 mol HCL
- 1 mol H2SO4
- 1 mol Acetic and Citric Acid
- 1 mol NaOH
- Indicator; Universal, Methyl Orange, Litmus, Bromothymol Blue, Phenolphthalein
Method:
1. Add 5 drops of each chemical into separate clean test tubes and place on test-tube rack
2. Add 2 drops of universal indicator into each test tube. then observe and record colour
3. Repeat steps I and 2 for each indicator and rinse test tubes between each use
4. Connect the pH probe to data logger, then remove probe from buffer solution.
5. Measure and record the pH of 5 drops of each chemical in clean test tubes.
For Part 2: Measure the pH for 20ml of 0.1 mol HCL, H2SO4, Acetic Acid, And citric acid using
the data logger, Calculate the H+ conc for each and record and compare results in a table
o H2SO4 had greatest H+ Conc (Diprotic Strong Acid) followed by HCl (monoprotic strong
acid) then Citric Acid (Triprotic Weak Acid) and Acetic Acid (Monoprotic weak acid)
o Sulfuric acid; pH = 0.7|HCl; pH = 1.2 | Citric acid; pH = 2.9 |Ethanoic acid; pH = 3.3
o Degree of ionization compared, Indicated H2SO4 strongest acid, Ethanoic Weakest
6. Repeat the entire investigation 5 times and average your data logger reading
Diagram
Discussion
- The experiment showed that the more concentrated the soln, The more its acidic/basic
nature. pH of water is neutral (pH=7), While H2SO4 acidic (pH>7)| NaOH is basic (pH<7)
ACCURACY:
- Very accurate, the use pH meters gave reading to 2 dp, Hence accurate comparison
between chemicals, Also indicated acidic and basic nature.
- Test-tube were rinsed and dried between each test to minimise cross contamination
VALIDITY:
- The experiment is valid, Only one variable was changed, Other Variables were controlled
o CONTROLLED: The number of drops of each chemical/Indicator, Size of test tube,
Temperature of environment
o INDEPENDENT: Chemical Used |DEPENDENT: pH/Acidic/Basic/ Neutral nature of it
RELIABILITY:
- The experimental results are reliable because it was repeated several times and
consistent results were achieved, It also accounted for theoretically expected results
Justifications
- A pH probe was used instead of a universal indicator => More accurate, 2 dp
- Various conc of the same solution were used to determine the affect of conc on pH
- The pH probe was placed in buffer solution before each test to rest the pH probe, To
prevent cross contamination and to give more accurate pH readings
Conclusion: HCl and H2SO4 showed acidic nature, NaOH showed basic while water neutral
Decarbonated Soft Drink Practical AIM: To observe the mass changes and carbon dioxide loss when soft drink is de-carbonated
by heat at 25oC and kPa.
EQUIPMENT
- Hot plate
- Electronic balance
- 250mL beaker
- Metal tongs
- 200mL of soft drink
SAFETY:
- Hot water or acidic soft drink may come into contact with your eyes which cause
discomfort, burning and long-term damage. To minimise these risks wear safety goggles.
- Hot plates are dangerous to touch and may cause burns, caution should be taken around
these and they should be allowed to cool before being handled.
METHOD
1. Pour 200ml of carbonate soft drink into a 250ml beaker
2. Weigh this on an electronic balance
3. Place the beaker on the hot plate at a medium heat setting
4. Re-weight the beaker every minute on electronic balance, use metal tongs to move it
o Repeat step 4 until consistent mass is achieved
5. Record your results in a table
Diagram
Results
Minute Mass (g) Minute Mass (g)
0 318.26 6 316.7
1 317.9 7 316.4
2 317.6 8 316.1
3 317.4 9 315.7
4 317.2 10 315.4
5 316.9
Calculations
DISCUSSION
- The equilibrium that exists in soft drinks is CO2(g) + H2O (l) H2CO3 (aq) ΔH < 0
- The experiment showed that as the soft drink was heated its total mass was decreasing,
o When heat is added to the system, it tries to counteract this change by favoring the
endothermic reaction. Hence equilibrium shift toward the left this produces more CO2
Validity
- Overall experiment is valid, Only one variable was changed, Other Variables controlled
- Control: volume of soft drink used, Time between each heating, Beaker size
- Independent: Time taken |Dependent: Amount of CO2 lost
Accuracy
- Fairly accurate as an electronic balance was used gave readings to 2 dp
- However we assumed that all mass lost was due to CO2 escaping vessel, but some water
vapor may have also escaped. To prevent this a control should have been used (a beaker
filled with distilled water) in order to account for this mass loss
- Also we assumed the experiment was carried out at 25OC for calculation, however this
temp could not be accurately maintained and results may have been affected by this
RELIABILITY:
- The experimental is reliable because it was compared with other member of the class
and consistent results were achieved, It also accounted for theoretically expected results
JUSTIFCIATION
- We used a hot plate instead of a Bunsen burner, as the heat form a hot plate is more
subtle and is a Bunsen burner was used a lot of water would have evaporated
Conclusion: The mass of the soft drink changed by X grams as the softdrink was
Decarbonated, and released Y L of CO2 at STP
Identifying the pH of Salt Solutions AIM: To determine the pH of various salt solutions
EQUIPMENT
- 5 ml 0.1 mol/L Salt solutions
- 5mL distilled water
- Test tubes + Rack
- pH Probe
- Data Logger
- Safety Goggles
METHOD:
1. Place 5mL of each salt solution in a separate beaker.
2. Take the pH probe out of buffer solution before placing it in the Salt solution.
3. Record the pH reading on the data logger.
4. Repeat steps 2 and 3 for all chemicals and water, rinse pH probe between measurement
5. Repeat entire experiment 10 times
Results
ACCURACY:
- Very accurate, The use pH meters gave reading to 2 dp, Hence accurate comparison
between chemicals, Also indicated acidic and basic nature.
- Test-tube were rinsed and dried between each test to minimise cross contamination
Validity
- Overall experiment is valid, Only one variable was changed, Other Variables controlled
- Control: Size of each test-tube, Volume of Solution used,
- Independent: Salt Solution |Dependent: pH of solution
RELIABILITY:
- The experimental is reliable because it was compared with other member of the class
and consistent results were achieved, It also accounted for theoretically expected results
Conclusion: Showed that salt solutions have a variety of pH
Preparing A Standard Solution AIM: To prepare a 1M standard solution of Na2CO3
EQUIPMENT
- Volumetric flask
- Sodium carbonate
- Electronic balance
- Glass funnel
- Pasteur pipette
- Distill water
- 200mL beaker
RISK ASSESMENT
- The solution produced will be slightly caustic, so gloves and safety glasses must be worn
METHOD
1. Weigh out 2.65 grams of Anhydrous sodium carbonate using an electronic balance into a
clean dry beaker. Add 50 ml of distilled water and stir with rod until completely dissolves.
2. Pour the solution into the cleaned volumetric flask (Rinsed distil water) using filter funnel
3. Wash out the beaker thoroughly and pour washings into volumetric flask.
4. Top up the volumetric flask with water until the bottom of the meniscus at 250mL mark.
5. Stopper flask, invert and swirl until solution becomes homogeneous.
DISCUSSION
The solution made may not have been accurately of concentration 1M because
- Equipment is calibrated at 25OC, and as this temperature could not be maintained, it may
cause causes errors when filling water to the graduation mark
- Water may have been added beyond or under the graduation mark, this could be
avoided by reducing parallax error and using a magnifying glass
JUSTIFICATION
- We used a volumetric flask instead of a conical flask, as a volumetric flask is specially
calibrated and accurately holds 250mL of solution
- The primary standard was dried in order to order to get rid of excess moisture.
- We used distilled water over normal tap water as it contained less impurities
- We also inverted the flask 5 times to ensure a homogenous mixture.
Titration Practical
AIM: To use standadised Na2CO3 and titrate it against HCI to determine its concentration.
Equipment
- Electronic Balance
- Small Beaker/Funnel
- Pipette (25ml) +Pipette Filler
- Burette
- Burette Clamp + Retort Stand
- 250 ml Conical flask
- Wash Bottle With Distilled water
- 0.1 mol/L HCl (100ml)
- White card
- Standardized 0.1M Na2CO3
- Methyl Red indicator
Method
1. Rinse pipette & burette with water then 5mL of HCl and sodium carbonate respectively.
2. Use a pipette to extract 25mL of HCI from the supplied container and pour it into a
conical bottom flask (Rinsed with water).
3. Add 3 drops of the indicator into flask and put a white paper underneath.
4. Fill a burette with a 0.100M solution of sodium carbonate.
5. Clamp the burette onto a retort stand and slowly drip sodium carbonate into the flask.
6. Continue stirring the flask then close the burette tap when the flask changes colour
7. Calculate the volume of sodium carbonate used and use it to find the conc. of HCl
8. Repeat the entire titration 5 times and average your results
Results
Calculations
ACCURACY/JUSTIFICATION:
- Anhydrous (normal) Na2CO3 was used because it is stable and does not gain mass by
absorbing water from air. Therefore, its composition and mass is accurately known
- Methyl Orange indicator was used, Appropriate changed from red to blue at 3.1- 4.4
which is close to the equivalence point of a strong acid and a weak base (< 7)
- All equipment was free of foreign contaminants and was washed with the correct liquids
- All volumes were measured from the bottom of the meniscus
- No splashing or loss of liquid occurred during the transfer of liquids.
- Hence accurate and valid as measurements were also close to stoichiometric values
Validity
- Overall experiment is valid, Only one variable was changed, Other Variables controlled
- Control: Size of pipettes/Burettes/Comical flask, Vol Solution used, # drops indicator
- Independent: Amount of Na2CO3 used|Dependent: Concentration of acid
RELIABILITY:
- The experimental is reliable because it was compared with other member of the class
and consistent results were achieved, It also accounted for theoretically expected results
Conclusion: A standard solution of sodium carbonate was successfully prepared and titrated
against HCl to find its concentration as 0.102 mol/L.
Titrating A Household Substance AIM: To find the concentration of acetic acid in household vinegar using tech (Data Logger)
EQUIPMENT:
Household vinegar
Distilled water
Burette and clamp
Glass funnel
Pipette and filler
Volumetric flask
Conical flask
Retort stand
White card
Data Logger + Comp
Phenolphthalein indicator
Standardized 1M NaOH
Method 1. Rinse pipette and burette with water and then 5mL of vinegar and NaOH respectively 2. Use a pipette to add 25mL sample of the vinegar into a 250mL volumetric flask. 3. Add distilled water to the fill the flask to the 250mL mark. 4. Use pipette to extract 25mL aliquot (sample) of diluted vinegar add this to conical flask. 5. Fill your burette with a 0.1M solution of NaOH. 6. Clamp the burette to a retort stand and position the volumetric flask under it. 7. Connect the pH probe to the computer and then remove the cap from the pH probe. 8. Place your pH probe into the vinegar flask, holding it upright through the experiment. 9. Open burette tap to let NaOH run into the flask I mL at a time until pH reaches
o Note reaction is between strong acid and strong basic so produce neutral salt.
10. Obtain a titration graph using your computer
Validity
- Overall experiment is valid, Only one variable was changed, Other Variables controlled
- Control: Size of pipettes/Burettes/Comical flask, Vol Solution used,
- Independent: Amount of NaOH used|Dependent: Concentration of acetic acid
Accuracy
- A pH probe/Data logger used and the computer recorded the results. This significantly
increased accuracy pH correct to 2dp gives a quantitative results unlike an indicator
- All equipment was used correctly and washed with correct techniques
- Experimental errors while reading the values on the pipette and burette, could be
minimized by reducing parallax error and using a magnifying glass
- The vinegar was diluted in order to ensure an accurate volume was used
- Improvement: We could have had a rough run to approximate when to slow the rate of
NaOH into flask and could have used indicator to confirm (validate) equivalence point
Conclusion
The concentration of acetic acid was found to be X and had an equivalence point of Y
PREPARING AN ESTER AIM: To prepare an ester using reflux 1-pentyl ethanoate (banana flavor)
EQUIPMENT:
Condenser
Round bottom flask
Heating mantle
Retort stand
Boiling chips
Conc sulfuric acid
1-pentanol
Ethanoic acid
Separating funnel
Distilled water
Sodium carbonate solution
RISK ASSESMENT
- Vapors produced by reaction are highly flammable so should be kept away from flame
- Substances used in experiment are highly corrosive so safety glasses/gloves worn
- Sulfuric acid is corrosive. Clean up spills immediately
METHOD
1. Place 10mL of 1-pentanol, 10mL of ethanoic acid, 1mL of sulfuric acid and a few boiling
chips into a round bottom flask., And place inside a heating mantle
2. Fix a water condenser upright into the flask using a retort stand
3. Turn on the water supply to condenser , And heat mixture under reflux for 30 min
4. Carefully pour the mixture into the separating funnel and add 10mL of water to it.
5. Stopper and shake funnel, allow layers to separated and discard the lower aqueous layer
6. Add 15mL of sodium carbonate solution, shake and discard the lower layer,
7. The layer remaining in the flask is the ester, which is further purified by distillation
Results: Initially ester smelt of nail polish before it was distilled, after it smelt like banana
DISCUSSION
- The ester was first washed with water to dissolve any remaining reagents and remove
this from the mixture, Since esters are insoluble (form an organic layer over the top. )
- The sodium carbonate was added in order to neutralize any remaining sulfuric acid
JUSTIFICATIONS
- Conc sulphuric acid used acts as a catalyst speeds up reaction, Dehydrating agent
- Boiling chips to promote even heating and prevent bumping
- Reflux was used to prevent the escape of any volatile vapors, allowing us to perform the
reaction at higher temperatures, boost the reaction rate and shift ↔ to favour ester yield
- We used an indirect heating source (not a Bunsen burner) to promote even heating and
prevent any vapours (Volatile Substance) produced igniting. Also increase safety
Conclusion: We were able to prepare the ester 1-pentyl ethanoate (banana flavor)
Deduce the ions present in a sample from the results of tests.
Aim: Measure the sulfate content of lawn fertiliser and explain the chemistry involve
EQUIPMENT
- 5g fertilizer
- Electronic balance
- Stirring rod
- 500ml beaker
- Sintered glass funnel
- 50 ml hydrochloric acid
- Bunsen burner
- Distilled water
- Mortar and pestle
- Barium chloride solution
RISK ASSESMENT
- HCl is highly corrosive and can burn the skin and eyes on contact, thus gloves and safety
glasses should be worn.
METHOD
1. Grind the fertiliser using a mortar and pestle into to a fine powder
2. Weight out 5g of fertilizer using an electronic balance, and carefully transfer all of it into a
500ml beaker, with the help of a wash bottle.
3. Add 100 mL of water and 50mL hydrochloric acid to the beaker and stir until all fertilizer
dissolves
4. Barium chloride was slowly added till no more precipitate formed
5. Heat the mixture below its boiling point (digesting) for 30mins and stir every now and then.
6. The sintered glass filter was weighed, and then the mixture was filtered through it.
7. The filter was then washed with distilled water and ethanol to remove any impurities
8. The filter was then dried in an oven until its mass became constant
9. The filter was re-weighed and the mass of barium Sulphates recorded
10. Calculate the mass of Sulfate and the % of Sulfate in the original sample of lawn fertiliser
Problem encountered Solution
Not all the fertilizer dissolved (if
this occurred not all sulfate would
dissolve leading to inaccurate
results) (some sulfate compounds
are insoluble)
We finely ground up the fertilizer to aid in dissolving
HCl to further aid dissolution
Some precipitate was lost during
filtration (the solubility of barium
sulfate is low so some may passes
through the filter un-dissolved)
We used a sintered glass filter as opposed to filter paper
as the larger pore size of filter paper would allow more
precipitate through
We allow the mixture to digest for 30mins which helped
to clump the precipitate together and made it easier to
filter
The mixture was cooled in ice water to reduce the
solubility of barium sulfate and cause it to precipitate out
Slowly forming precipitates using a dropper bottle in
high temperatures promoted the growth of larger
clumps of precipitate
Incomplete drying of the glass filter The filter was dried in an oven
Could be improved by placing in a desiccator
Fertilizer remained in mortar,
precipitate still stuck to beaker
The equipment was washed into the next stage to
remove any traces of remaining particles.
Heterogeneous composition of the
fertilizer
The experiment is repeated 5 times and outliers
excluded from averages.
Aim: Perform a first-hand investigation to model an equilibrium reaction
EQUIPMENT
- 2 x 100mL measuring cylinder
- 10mL pipette
- 2mL pipette
- Water
METHOD
1. One 100mL measuring cylinder (cylinder 1) was filled with 100mL water while the other
one was left empty (cylinder 2)
2. A 10ml pipette was placed into cylinder 1, and the water was allowed to rise, this volume
was then transferred to cylinder 2 [Represents Forward Reaction: Water A Water B]
3. The 2mL pipette was then placed in cylinder 2, and the water was allowed to rise, this
volume was then transferred to cylinder 1 [Backwards reaction: Water B Water A]
4. The volume of both cylinders was recorded
5. Steps 2 and 3 are repeated for 20 times
DISCUSSION
- Cylinder 1 represents the concentration of reactants, and the volume transferred from it
refers to the forward reaction
- Cylinder 2 represents the concentration of products, and the volume transferred form it
refers to the reverse reaction
- When the depth of water is high in a cylinder (i.e. the concentration is high) the volume
transfers increases, which is consistent with Le-Chatelier’s principle which suggest hat if
the concentration of a species is high, the equilibrium shifts towards the other side of the
equilibrium
- When the volumes becomes constant, the rate of forward reaction has reached the rate
of reverse reaction and our system is in equilibrium
- Equilibrium could be disturbed by adding extra water into Measuring cylinder A
Aim: To qualitatively analyses an equilibrium reaction.
- Iron(III) ions react with colourless thiocyanate ions, SCN–
, to form an intensely red-
coloured complex in an equilibrium reaction: Fe3+
(aq) + SCN– (aq) Fe(SCN)
2+(aq)
Method
- Using a clean pipette for each, transfer 5ml of 0.1 mol/L iron (III) nitrate Fe(NO3)3 & 5 ml
of 0.1 mol/L potassium thiocyanate KSCN solution respectfully, into clean 250 mL beaker.
- Add 100 mL of distilled water and mix the chemicals thoroughly
- Add 10 mL of this mixture was added to each of the 4 medium size test tubes.
- Tube 1 was a control; It enabling colour (And intensity) changes to be compared to.
- The other test tubes were treated as follows
o Tube 2: Added 10–15 drops of Fe(NO3)3 solution
o Tube 3: Added 10–15 drops of KSCN solution.
o Tube 4: Added a few drops of 2 mol/L NaOH solution.
o Tube 5: Add 1.0 g solid potassium nitrate.
o Tube 6: Heated in a water bath.
o Tube 7: Placed in an ice cold water bath.
- Colour changes and colour intensity changes were observed with reference to the
control, and results were recorded in a table.
Results: Explain how you analysed the equilibrium reaction qualitatively.
- Addition of NaOH (aq) results in the OH- reacting with the Fe3+ ions forming a ppt of
Fe(OH)3. Conc Fe3+ decreases by LCP backward reaction favoured, Lighter red colour
- Adding KSCN solution results in the conc of SCN- to increase, According to LCP the
forward reaction rate is favoured, Red Darker red colour.
- Heating the mixture resulted in the backwards endothermic reaction to be favoured by
LCP, Blood Red Lighter Red (decrease to the colour intensity)
- Cooling the mixture resulted in the forward exothermic reaction to be favoured by LCP to
remove heat, Blood Red Darker Blood red (decrease to the colour intensity)
Light Yellow Colourless Blood Red
Aim: To identify the products of the electrolysis of sodium chloride
RESULTS/Discussion:
- In the dilute solution we observed bubbles forming at the 2 electrodes indicating gasses
o Oxidation Half equation (Anode) – H2O (l) 0.5O2(g) + 2H+ + 2e-
o Reduction half equation (cathode) – H2O (l) + e- 0.5H2 (g) + OH-
(aq)
- When a glowing splint was placed over the anode compartment, the splint re-ignited this
confirmed the presence of oxygen gas. When a burning taper was placed over the
cathode a pop sounds was heard, this confirmed the presence of hydrogen gas.
- In the concentrated solution we also witnessed bubbles forming at the 2 electrodes
o Oxidation half equation (anode) – Cl- 0.5Cl2 (g) + e-
o Reduction half equation (cathode) – H2O (l) + e- 0.5H2 (g) + OH-
(aq)
- Litmus paper placed over the anode it was seen to be bleached, indicates chlorine gas
- Match over the cathode made a pop sounds, Indicates hydrogen gas.
- When phenolphthalein indicator was added a dark pink color was produced indicating
the presence of OH- ions, as its changes color from colorless to pink in basic conditions.
Aim: To prepare soap by saponification of a fat and to test the properties of the soap.
Equipment:
- 10 g NaOH pellets
- 10 g Coconut oil
- Brine (NaCl solution)
- 500ml large beakers
- 40ml water
- Bunsen burner+Tripod+Gauze mat
- Commericial soap
- 4 testtubes
Risk Assessment:
- Safety glasses and gloves were worn, as sodium hydroxide is caustic and must not be
splashed onto the skin or into the eyes. The experiment was also carried out near a tap
so that a supply of running water was available.
- The soap was washed washed before handling to remove any traces NaOH
Method:
1. Measure about 10 g of NaOH using balance, Dissolve the NaOH in about 40mL of distilled
water in a clean 500mL beaker.
2. Measure and add 10 g of coconut oilto beaker.
3. Heat the apparatus using a Bunsen burner for approx. 30mins while constantly stirring.
Maintain the volume by adding a little distilled water. If the reaction has finished the oily
layer should have disappeared. The mixture usually looks thickened and soapy.
4. When the mixture has cooled, pour it into a solution of brine (NaCl solution) to
precipitate the soap out as hard lumps or 'curds'
5. Decant the solution into a waste container. Wash the soap with NaCl solution twice,
decanting the solution into waste container each time.
6. Tip the soap onto paper towel to dry.
PART B: Testing the soap
1. Place 3 grams of the soap into two clean test tube, and place the same amount of
commercial soap in two test tubes.
2. Add about 3 mL of distilled water into the first of the pair of test tubes and add 3 mL of
hardened water to the second of each pair of test tubes.
3. Shake each test tube vigorously and record height and density of foam produced
Results:
- Our soap acted like a commercial soap but not as well because it did not foam to same
height or with the same density as the commercial soap because these contain extra
additives that make them foam better
DISSCUSION: Coconut oil was used as it cheap, pure and readily available,
- We used a water to heat the reagent gently
Sulfuric acid as a dehydration/oxidizing agent
- AIM: To observe sulfuric acid acting as a dehydrating and oxidizing agent
EQUIPMENT:
- Sucrose (sugar)
- Concentrated H2SO4
- Copper metal
RISK ASSESMENT:
- Fumes releases form conc. H2SO4 can be toxic, to the experiments are performed in a
fume cupboard
- Safety goggles and gloves must we worn at all time as H2SO4 is highly corrosive
METHOD:
1. Place a strip of copper metal into a 10ml solution of concentrated sulfuric acid
2. Record observations
3. Place 10g of sucrose into a 800ml beaker
4. Add 10ml of concentrated H2SO4, and observe the reaction.
OBSERVATIONS/RESULTS:
- During the oxidizing experiment, we observed the solution became increasingly blue in
color, we also observed the copper metal strip reducing in size, and a gas being released.
- During the dehydration experiment, we observed a gas escaping a large black tower
forming.
DISCUSSION
- During the oxidization practical we observed the following chemical reaction:
- 2H2SO4 (aq) + Cu(s) SO2 (g) + CuSO4 (aq) + 2H2O (l) (*Note: SO2 and H2O forms when
CONCENTARTED sulfuric acid reacts with metal.)
- The blue color we observed was attributed to the increasing concentration of copper
ions in the solution, making it blue. The gas released was toxic sulfur dioxide, and the
reducing size of the copper metal was caused by the conversion of solid copper to ions.
In the equation above we see the oxidation state of copper change from 0 to 2+,
therefore it is being oxidized, and hence H2SO4 is acting as an oxidizing agent.
- During the dehydration experiment we observed the following chemical reaction:
- C12H22O11 11H2O(l) + 12C(s)
- The gas escaping was a result of water vapor forming form the sucrose, the large black
tower, was made of carbon produced during the dehydration.
Conc. H2SO4
JUSTIFICATIONS:
- We used cooper as our chosen metal as the formation of blue solution makes it clear the
copper ions have been formed, and thus copper has been oxidized
- We also used sucrose to dehydrate, rather than other substances such as glucose, as this
reaction is clearly visible due to black tower of carbon forming, and the steam produced
makes it clear that dehydration has taken place.
CONCLUSION: we were successfully able to observe sulfuric acid acting as a dehydrating
agent and an oxidizing agent.
Perform a first-hand investigation to demonstrate the effect of soap as an emulsifier
1. Add 100 mL of water and 5mL of oil to a flask and stopper it. Shake the flask and then
leave it to settle, observing the formation of any layers.
2. Remove the stopper and add about 5 mL of liquid soap, Shake the flask and then
leave it to settle, observing the formation of any layers.
- Separate layers of oil and water formed as the mixture is the flask settled after step 1.
This is because the non-polar oil does not dissolve in the polar water (they are immiscible)
- When the liquid soap was added and the flask shaken, the separate layers disappeared
and the mixture became homogenous, although a foam formed on the water's surface.
This is because the soap anions were able to 'solubilise' the oil in the water by forming
micelles with the oil due to the soap anion's structure (hydrophobic tail and hydrophilic
head). This enabled the oil to become evenly and indefinitely dispersed through the
water, forming an emulsion. Hence, the soap acted as an emulsifier
Perform a first-hand investigation to gather information and describe the properties of a
named emulsion and relate these properties to its uses:
Mayonnaise:
- Mayonnaise is an emulsion of vegetable oil and egg yolks, with the emulsifier being the
lecithin found naturally in the egg yolk.
- Other additions may be made for flavour, such as vinegar, mustard of salt,
- Mayonnaise is made by slowly adding oil to an egg yolk, while whisking vigorously to
disperse the oil; the lecithin stabilises the mixture.
Properties in Relation to Uses:
- Mayonnaise is a very STABLE emulsion, due to lecithin. It does not separate into its
component liquids even when stored for long periods of time, useful food product edible
for longer. Has a creamy ‘mouth-feel’, and not feeling oily enjoyable taste.
- However, mayonnaise is actually on average 75% oil. The property of the emulsion as
having a creamy texture adds to its use as a food.
- AIM: To model the formation of sodium bicarbonate in the Solvay process
- RISK ASSESMENT:
The ammonia solution and ammonium crystals produce ammonia gas, which is highly
toxic even in small quantities, may lead to unconsciousness and asthma this risk is
minimized my performing the experiment in a fume cupboard
The dry ice can be -80OC, this is cold enough to cause skin irritant and burns to skin,
thus dry ice should only be handled with gloves and tongs
Prevent ammonia and dry ice from burning the skin and eyes, wear safety goggles and
non-rubber, insulating gloves.
Method
1. In a conical flask ammonia solution was added to 20ml of brine until completely
saturated, this will ammoniate the brine.
2. Add 2 cubes of dry ice using tongs.
3. Observe and record formation of any precipitate. A white precipitate should be
observed, if no precipitate form cool the apparatus.
4. To complete the modelling of the Solvay process, we would need to filter out the
NaHCO3(s) from the solution and decompose it to form sodium carbonate by heating
Discussion
- Dry ice was used to supply carbon dioxide for the reaction to occur and to reduce the
temperature of the solution toabout OoC to enable the precipitation of NaHCO3 (it only
does this at OoC).
- The reaction that occurred was:
- NH3(aq) + CO2(g) + NaCl(aq) )NH4Cl(aq) + NaHCO3(s)
- The NaHCO3(s) precipitated out.
Difficulties:
- It was difficult to crystallize out the NaHCO3(s) the mixture could not be cooled
homogenously. It was also difficult to recover the ammonia because it is so hazardous
and further steps would be required, thus, it was avoided.
- In the Solvay process the CO2 is sourced from the decomposition of limestone (CaCO3 ),
the lab lacks equipment to sufficiently heat limestone to thermally decompose it , hence
dry ice is used