(Topic 1) Chemistry Lab – Finding the Formula of a
Hydrate
Purpose:
To find the molecular formula of a certain hydrate of copper (II) sulfate,
CuSO4 • xH2O. You will be working in groups of 3-4. However, each
person will submit their own lab report.
Assessment:
CE – record and label your results carefully, including units, significant
figures, and uncertainties. You will be comparing this to the true value
of the hydrate. Also pay attention to any variables that should be
controlled during this experiment.
Materials:
Ceramic evaporating dish
Electronic balance
Bunsen burner
Retort stand, ring clamp, clay triangle
Glass stirring rod, scoopula, crucible tongs
Copper (II) sulfate hydrate crystals
Safety:
 All students must wear safety glasses at all times during the lab,
even when not working directly with the chemicals or apparatus.
 Be careful around the hot plates, even when they are unplugged
– they may still be hot!
 Do not inhale the copper (II) sulfate hydrate
 Do not touch the copper (II) sulfate crystals
 Always wash your hands thoroughly before and after handling any
chemicals.
Procedure:
1. Gather all the equipment you need for this lab.
2. Use evaporation to determine the formula of CuSO4 • xH2O. Use
between 3-5g of crystals per trial.
3. Place the clay triangle on top of the ring clamp, and set the
evaporating dish on top.
4. Light the Bunsen burner. You will hold the Bunsen burner
throughout the entire experiment and wave it gently back and
forth underneath the evaporating dish.
5. Use your stirring rod to CAREFULLY mix the crystals in the dish so
that they do not burn. Continue allowing the crystals to heat
until ALL the crystals have changed colour. Do not lose any
crystals during the heating. Record any observations you notice
during the heating of the crystals.
6. If you hear sizzling, then the heat is a little bit too high.
7. Once the colour change has finished, use the tongs to remove
the evaporating dish from the clay triangle. Be careful, it is very
hot.
8. Let the sample cool down for 1-2 minutes before measuring.
9. Place the anhydride into the storage beaker provided by your
teacher. Wipe the evaporating dish using paper towels.
10. Repeat all steps of the experiment a total of 3-4 times,
depending on time.
Analysis/Calculations:
1. Find the average of your 3-4 trials.
2. Use your average values to calculate the molecular formula of
the hydrate (the value of x).
Conclusion:
What is the molecular formula of this hydrate?
++++++++++++++++++++NOTES+++++++++++++++++++++
Find the formula of CuSO4 5H2O from experiment.
CE
Don’t give them too much information about what to measure.
Save the CuSO4 anhydride crystals for resuse, or to make simple CuSO4 solutions.
(Topic 1) Empirical Formula Lab: Magnesium Oxide
Introduction:
Metal oxides are formed when a metal is reacted with
oxygen. In this lab, you will convert Magnesium metal (an
element) to magnesium oxide (a compound). Magnesium is an
alkaline earth metal, which reacts vigorously when heated in
the presence of air to produce magnesium oxide and magnesium
nitride. We only want magnesium oxide, however, so water is
added to react with the magnesium nitride and convert it to
magnesium hydroxide. This is then further oxidized by burning
to convert all of the Mg to magnesium oxide. Based on the
weights of the reactants and the product, the percent
composition of the product can be found. Using this
information, the empirical formula of the oxide, MgxOy can be
determined.
Purpose:
To discover the empirical formula of magnesium oxide by
oxidizing magnesium.
Materials:
Magnesium ribbon
Distilled water
Ceramic crucible & cover
Bunsen Burner
Ring stand
Clay triangle
wire gauze
Crucible tongs
Electronic Balance (at least 0.01g accuracy)
Safety:
Burning magnesium produces a very bright white flame. DO NOT
look directly at the flame, it can damage your eyes.
Crucibles can become very hot but still look cold. Be careful
around any crucibles because they may be extremely hot.
Always use crucible tongs to carry crucibles carefully, this
is also so that oils and impurities from your body do not get
onto them.
Procedure:
1. Place a clean, dry crucible and crucible cover on a clay
triangle on a ring stand over a Bunsen burner. Adjust
the height of the ring so the bottom of the crucible will
be in the hot part of the flame. Place the cover so it is
slightly ajar, allowing air to leave the crucible.
2. Heat the empty covered crucible for about 10 minutes. The
bottom of the crucible should be red hot. Turn off the
burner and allow the crucible to cool for 5 minutes.
3. When the crucible is cooled to room temperature, transfer
it to the balance with crucible tongs and record the
mass.
4. Prepare to heat the magnesium inside the crucible.
Record whatever measurements you need. Never touch the
crucible with your hands or place on the lab bench, this
may result in contamination.
5. Obtain about 0.15-0.2g of magnesium ribbon. If the metal
is not shiny, rub it with steel wool to remove oxidized
metal.
6. Fold the magnesium ribbon and then place it inside the
crucible.
Place the crucible on the clay triangle over
the Bunsen burner and start heating.
7. As soon as the magnesium glows, place the cover on the
crucible, so that no smoke escapes. Periodically lift
the cover to check the progress of the reaction and to
allow oxygen into the crucible. When the reaction is
complete, remove from heat and place on wire gauze.
8. Add 10 drops of distilled water and observe, including
smell.
9. Heat the crucible slightly covered until the water has
completely evaporated. Allow the crucible to cool down
for 5 minutes, then record readings.
10.
Repeat the procedure 2-3 times.
Calculations:
1. Determine the mass percent composition of your product.
What are we assuming about the conversion during this
reaction?
2. Determine the empirical formula of magnesium oxide.
3. Find out the correct empirical formula of magnesium oxide
and compare your results. Give some possible reasons for
any discrepancies.
4. Why is it important to heat the empty crucible at the
beginning of the lab?
5. Write a balanced chemical equation for the conversion of
magnesium to magnesium oxide, using the experimental
empirical formula found for MgxOy.
Conclusion:
Write a suitable conclusion for this lab.
(Topic 1) Chemistry Lab – Limiting and Excess Reactants
Introduction:
To investigate the reaction between aluminum and aqueous copper (II)
chloride. The blue colour of the CuCl2 solution is caused by the Cu2+ ion. Here
is the balanced chemical equation:
2 Al(s) + 3 CuCl2(aq)  3 Cu(s) + 2 AlCl3(aq)
Purpose:
To determine the concentration of an unknown CuCl2 solution. You will react
with varying quantities of aluminum and try to determine the concentration of
the CuCl2 by stoichiometric calculation of solutions.
Materials:
100mL Beakers, Graduated Cylinder
Electronic balance
Glass stirring rod, scoopula
Aluminum foil
Copper (II) chloride solution (unknown concentration)
Procedure:
11. Gather all the equipment you need for this lab.
12. Use a graduated cylinder to pour 50.0mL of the CuCl2 solution into
each beaker.
13. Measure out some Al foil. It can be made of more than 1 piece. (hint:
the correct amount you need is somewhere between 0.15g and 0.25g.
14. Stir the beaker and observe the reaction. When the reaction is
complete, finish your observations, choose a different quantity of Al,
then go back to Step 2.
15. Carefully pour the beaker contents into the aqueous waste jar.
16. Wash all the equipment you used and return them to their proper
places.
Analysis/Calculations:
3. How can you tell when CuCl2 solution is in excess?
4. How can you tell when Al foil is in excess?
5. One of your trials should have CuCl2 in excess. Calculate the amount
of each substance that reacted in this trial.
6. Determine the concentration of the CuCl2 solution to as many
significant figures as possible.
7. If the concentration of CuCl2 was 0.1500M, how many grams of CuCl2
solid would be required to make exactly 1.000dm3 of solution?
8. If the concentration of CuCl2 was 0.1800M, how many grams of
CuCl22H2O hydrate solid would be required to make exactly 750 cm3
of solution?
++++++++++++++++++++++++NOTES++++++++++++++++++++++
Refer to p255 in McGraw-Hill Ryerson Chemistry 11 textbook for template.
Adjust the lab slightly for the class
CuCl2 + Al foil lab.
CuCl2 concentration:
700mL of 0.25M diluted to 1.000L = 0.175M
For 50mL requires about 0.1575g of Al foil.
You can simply add slightly more and more Al each time and stop when the last Al
does NOT react anymore! Unless you use too much of course
(Topic 1) Quantitative Chemistry Lab: Deduce the HCl
Concentration
Introduction:
HCl reacts with Mg metal to produce H2 gas in a slightly exothermic
reaction. Using the balanced chemical equation and basic
stoichiometry, we can calculate the concentration of an unknown HCl
solution. We can measure the amount of H2 gas that was produced in
the reaction by collecting it in an upside down graduated cylinder.
The concentration of the HCl is somewhere between 0.5M and 1.0M.
Purpose:
To quantitatively determine the concentration of an unknown HCl
solution using the reaction between Mg and HCl that produces
hydrogen gas.
Assessment:
DCP and CE
Think about what data you will need and record it carefully, including
units, significant figures, and uncertainties.
Materials:
Mg ribbon
HCl (unknown concentration)
Erlenmeyer flask with gas delivery nozzle
Rubber tubing
Rubber bungs
100mL Graduated cylinders
10mL Pipettes and Pipettors
Large Plastic Tray
Diagram:
Figure 1: Gas delivery setup - tubing inserted in inverted graduated cylinder
Procedure:
1. Gather the materials listed
2. Fill up the graduated cylinder with water completely, then carefully invert it
inside the tray of water so that no gas bubbles are inside the upside-down
cylinder.
3. Setup the gas delivery tube so that it feeds directly into the cylinder, so that
none of the gas will escape, and connect the other end to the Erlenmeyer
flask gas delivery nozzle.
4. Obtain a sample of acid (about 10mL) and place it in an Erlenmeyer flask.
How are you going to collect this sample to as many significant figures as
possible?
5. Obtain at least 8cm of Mg ribbon.
6. When you are ready with the stopwatch and the tubing, place the Mg inside
the flask. Immediately start timing, cover the flask with the bung and hold
the tubing inside the cylinder. 2 people should be operating the apparatus
and 1 person should be recording data.
7. Take volume measurements every 30 seconds until the reaction finishes.
8. The final volume of gas is not only hydrogen gas, it has water vapour as well,
because the gas was bubbled through water. You must include this water
vapour pressure in your calculation.
9. Repeat this process 2 more times. Use your results to calculate the
concentration of the HCl solution.
10. Test the gas produced with a flaming splint.
11. Used acid should be thrown in aqueous waste, and the magnesium strips
should be collected and returned to your teacher.
Observations:
Record all necessary observations including units, significant figures, and
uncertainties in an appropriate format. Also pay attention to any possible sources
of error for CE.
Calculations:
Calculate the concentration of the HCl acid as accurately as possible. Present the
results in a suitable format.
Conclusion:
What is the concentration of the HCl acid? Evaluate the procedure based on its
validity and the controlling of relevant variables.
+++++++++++++++++++++NOTES+++++++++++++++++++++++
Measure:
Rate: time every 10 seconds to plot the rate of reaction
Quantity of H2 gas: obtain moles, calculate concentration of unknown HCl
Use Mg, way faster (Zn takes forever) , no less than 1M acid, around 20mL?
Suggested:
10mL of 0.800M HCl, 8cm strip of Mg, 10 minutes, 75-95mL gas
(Topic 5) Calorimetry – Enthalpy of Reaction
Introduction:
The reaction between copper (II) sulfate and zinc
metal is exothermic, and will produce heat.
The
amount of heat produced can be measured by
calorimetry.
This is a process by which the heat is
captured by a vessel (a calorimeter) and a medium
(usually water).
The amount of heat captured is
measured in terms of a temperature increase, and is
equal to the heat generated by the reaction.
Purpose:
To determine quantitatively the enthalpy of reaction
between copper (II) sulfate and zinc metal.
Assessment:
DCP and CE (but we will do this one together)
Think carefully about what measurements you need to
make.
Always consider uncertainties and significant
figured with numerical data.
Materials:
Copper (II) sulfate solution (1.00M  0.01M)
Zinc metal
Copper Calorimeter
Data logging software
Vernier Temperature probe
Weighing boat
Pipette and pipettor
Procedure:
1. Gather materials for lab.
2. Use a pipette to measure exactly 25cm3 of CuSO4
solution into the calorimeter.
3. Close the lid of the calorimeter and begin the
data logging software, set to a reasonable rate
of readings.
Wait until they stabilize.
4. Measure out about 2 grams of Zn powder into a
weighing boat.
This is the excess reactant.
5. Put all of the Zn metal in at once.
6. When the reaction has completed, continue
measuring the temperature for a few minutes to
let the temperature cool down a bit.
7. Dispose of the materials in the inorganic waste
beaker and wash out the equipment carefully.
8. Repeat this experiment a total of 3 times.
Data Collection and Processing:
Since we are using data logging software, you must be
careful to take the raw data and present it in your
own words and your own table.
Then you must process
your own data yourself, and not just use the
computer-generated graph.
Make sure your data tables, diagrams, and/or graphs
all have meaningful titles, units, significant
figures, and uncertainties, where appropriate.
Do not draw a line of best fit.
Extrapolate the graph data points backwards until the
appropriate time.
Why is this necessary?
Calculate the enthalpy changes in this experiment,
making assumptions
Calculate the enthalpy of the reaction between copper
(II) sulfate and zinc.
Conclusion and Evaluation:
Compare your value with the true value of H = -217
kJmol-1
State possible sources of error in this experiment
Describe ways these errors could be eliminated or
reduced
(Topic 5) Calorimetry – Hess’s Law, MgSO4 Hydrate
and Anhydrate
Introduction:
Hess’s Law is used to relate several different reactions
and their enthalpies.
The enthalpy difference from 1
state to another is the same regardless of how that
change was achieved.
Some changes cannot occur, such as
the direct hydration of MgSO4(s) to form MgSO47H2O(s).
Using Hess’s Law, however, we can measure separate
enthalpy changes involving both of those 2 states and
relate them using a 3rd state which is common to both
substances, in this case, their aqueous solutions.
Purpose:
To calculate the enthalpy of reaction for MgSO4(s) + 7
H2O(l)  MgSO47H2O(s)
Assessment:
DCP and CE
Think carefully about what measurements you need to make.
Always consider uncertainties and significant figured
with numerical data.
Materials:
MgSO47H2O(s) crystals
MgSO4(s) crystals
Copper calorimeter
Distilled water
Electronic scale
Data logging software
Vernier temperature probe
Stopwatch
Weighing boats
Procedure:
1. Calculate the appropriate quantities of each
compound that you need to add.
The number of moles
of MgSO4 must be the same in each case, and the
number of moles of H2O must also be the same.
You
must also make sure that those amounts will fit into
the copper calorimeter.
2. Gather required materials for the lab, and set up
the data logging software with the temperature
probe.
3. Take any initial measurements that you need to make.
4. Obtain the calculated quantity of water and place it
within the calorimeter.
5. Obtain the calculated quantity of MgSO4 crystals in
a weighing boat.
6. When you are ready, pour the crystals into the
calorimeter, record readings until the experiment
has finished, then take final measurements.
7. Repeat steps 3-6.
8. Repeat steps 3-7 with the MgSO47H2O(s) crystals.
Data Collection and Processing:
Make sure your data tables, diagrams, and/or graphs all
have meaningful titles, units, significant figures, and
uncertainties, where appropriate.
Calculate the enthalpy change of the hydration of MgSO4
crystals: MgSO4(s) + 7 H2O(l)  MgSO47H2O(s).
Explain why it is not necessary in this lab to
extrapolate the temperature time graph backwards.
Conclusion and Evaluation:
Compare your value with the true value of H = -104
kJmol-1
State possible sources of error in this experiment
Describe ways these errors could be eliminated or reduced
(Topic 5)Thermometric Titration: Acid –
Base Neutralization
Introduction:
Titration is a method by which a specific quantity is
neutralized by another substance and the exact
amounts are measured, usually using a burette.
The
end-point of the reaction is when 1 reactant has been
exactly neutralized by the other, in the
stoichiometric ratio.
Usually, the purpose is to
determine the concentration of an unknown liquid.
In
this case, you will also calculate the enthalpy of
the acid-base neutralization reaction.
Temperature
will be used to monitor the end-point, this is called
thermometric titration.
Purpose:
To determine the concentration of an unknown acid and
the enthalpy of neutralization for 2 acid-base
reactions, between NaOH and CH3COOH and HCl.
Assessment:
DCP and CE
Think carefully about what measurements you need to
make.
Always consider uncertainties and significant
figures with numerical data.
Materials:
Burette
Pipette and Pipettor
Data Logging Software and Vernier Temperature Probe
Standard base solution (NaOH 1.02M  0.01M)
HCl (unknown concentration)
CH3COOH (unknown concentration)
Copper Calorimeter
Funnel
Procedure:
1. Obtain exactly 25.00mL of the NaOH solution into
the calorimeter.
Insert the temperature probe
and initiate the software.
2. Clean the burette with water, then rinse once
with HCl acid.
3. Fill the burette to the top with HCl.
4. You are about to begin this exothermic reaction
and the thermometric titration.
In order to be
as accurate as possible and not lose too much
heat, you will have to work quickly but
accurately.
You must think carefully how you
are going to take readings.
5. Perform the thermometric titration and make sure
you go past the end-point so that the
temperature starts to decrease.
6. When you finish, repeat steps 1-4.
7. Repeat steps 1-5 using the CH3COOH instead of
HCl.
Data Collection and Processing:
Make sure your data tables, diagrams, and/or graphs
all have meaningful titles, units, significant
figures, and uncertainties, where appropriate.
Conclusion and Evaluation:
The enthalpy of neutralization for these 2 reactions
are given here.
These were obtained by the teacher
having done the experiment earlier.
H neutralization of HCl with 1.02M NaOH:
-49.26
kJmol-1
H neutralization of CH3COOH with 1.02M NaOH:
-52.83
kJmol-1
Why do you think there are no standard values for
enthalpy of neutralization of HCl and NaOH?
++++++++++++ NOTES for all 3 Energetics Labs +++++++++++++++
3 major DCP and CE labs here, also ICT using probes, graphing software
simulations??
Zn + CuSO4 DH = -216 kJ/mol
Very exothermic
With these specifications: 20-50°C in about 8 minutes, total 12 with lead and lag time.
Very easy to measure the max temperature reached
MgSO4: not very exothermic at all, need a sensitive measuring device
10 minutes to dissolve completely, requires lots of stirring
how will this be measured accurately?
Acid-base: somewhat exothermic, need a fairly accurate measuring device, or a fairly
high concentration
25mL 1M NaOH titrated with 2M acid takes about 12 mL and 6-7 minutes
turnaround, rise 6-8 °C
Look at Excel file: thermometric titration for results, plotting the graph, extrapolation,
etc.
Mass beforehand and after of calorimeter and stirrer and liquid
Density of liquid
Endpoint of titration: via graph = volume and max temp. (proportion of volume added
at endpoint
Q=mCDT to find heat released
Heat released goes into 0.025moles of the acid = enthalpy of neutralization
Or perhaps doing by rate would work. Set the probe to record every second, and get
ready to start the stopwatch or simply read time 0 from the graph. Set the burette to
drip every second, and just continuously stir the mixture. Should see the curve
exactly, and then when it declines, stop the burette flow and the stopwatch, calculate
the rate. Problem = rate changes as the level decreases because of pgh.
(Topic 6) Kinetics Lab – Factors Affecting Rate of Reaction
Introduction:
There are several factors that affect the rate of chemical reactions.
Two of the most important are concentration and temperature. In
this lab, we will investigate and quantify the effect of concentration
and temperature on the rate of a chemical reaction. This will be
measured in 2 different ways: through mass decrease, and
through volume increase. We will be using the reaction between
an acid and calcium carbonate:
CaCO3(s) + 2HCl(aq)

CaCl2(aq) + CO2(g) + H2O(l)
Purpose:
How does the temperature and concentration of the acid solution
affect the rate of reaction between calcium carbonate and
hydrochloric acid?
Assessment:
Make sure you collect appropriate data and process it accordingly.
Do not forget about units, uncertainties, and significant figures.
- Construct graphs of every trial
- Use the graphs to find initial rates of reaction
- Graph initial rate of reaction vs. Concentration
- Graph initial rate of reaction vs. Temperature
- Qualitatively describe the effect of changing temperature and
HCl concentration on rate of reaction
- Determine order of reaction with respect to HCl (HL
students)
- Determine the effect of temperature on rate of reaction (HL
students)
Part A: Concentration (measured by volume increase)
Controlled variables: temperature, volume of acid, mass of calcium
carbonate, measuring device, gas syringe
Independent variable: concentration of hydrochloric acid
Dependent variable: rate of reaction (rate of gas produced)
Procedure:
You will be measuring the rate of the reaction by measuring the
rate at which gas is produced (by feeding it directly into a gas
syringe).
1. Gather 1.0g of CaCO3 powder into a 250mL Erlenmeyer flask
with a gas discharge outlet. Have a bung ready to cover up the
top of the flask.
2. Attach a tube from the gas discharge outlet to an empty gas
syringe with the plug all the way in.
3. Obtain 50cm3 of 0.02 moldm-3 HCl in a 100mL graduated
cylinder.
4. Have a stopwatch ready and decide how often you will take
measurements.
5. When you are ready, pour the acid in, cover up the flask with the
bung, and begin the timer.
6. Stir the flask continuously throughout the reaction.
7. Take readings until the reaction is finished.
8. Dispose of the finished materials and clean the flask out as
much as possible.
9. Repeat the experiment for the following concentrations of acid:
0.05 moldm-3
0.10 moldm-3
0.20 moldm-3
0.50 moldm-3 (this experiment requires 2.0g of CaCO3)
Part B: Temperature (measured by mass decrease)
Controlled variables: concentration/volume of acid, mass of
calcium carbonate, measuring device, electronic balance
Independent variable: temperature of HCl
Dependent variable: rate of reaction (rate of mass decrease)
Procedure:
You will be measuring the rate of the reaction by measuring the
rate at which the mass of the reaction decreases.
1. Gather 1.0g of CaCO3 powder into a 250mL beaker.
2. Obtain more than 250mL of 0.1 moldm-3 HCl in a beaker and
begin heating it.
3. Once the temperature reaches 95°C, stop heating.
4. Use a graduated cylinder to measure 50cm3 of it into another
beaker.
5. Have a stopwatch ready and decide how often you will take
measurements. Place the CaCO3 beaker onto the electronic
scale.
6. As soon as the temperature drops to 90°C, pour the acid into
the beaker with the CaCO3, and start the timer.
7. Stir the flask continuously throughout the reaction.
8. Take readings until the reaction is finished.
9. Dispose of the finished materials down the drain with water.
10. As the acid cools down, repeat the experiment for the following
temperatures:
70°C
50°C
30°C
10°C (you will need ice to cool the acid)
Calculations (DCP or just course?)
Volume/Time:
1. Construct several plots of C vs. t
2. Find the initial slope of the graph at t=0, this will be in units cm3t-1
3. Change the units using PV=nRT to molt-1 (this is rate)
4. Each of your plots will ield 1 rate value
5. Using these rate values, plot 1 master Rate vs. Concentration graph
6. This graph represents the relationship between rate of reaction and acid
Mass/Time:
1. Construct several plots of m vs. t
2. Find the initial slope of the graph at t=0, this will be in units gt-1
3. Change mass to moles using M=m/n to molt-1 (this is rate)
4. Each of your plots will yield 1 rate value
5. Using these rate values, plot 1 master Rate vs. Concentration graph
6. This graph represents the relationship between rate of reaction and acid
concentration
(Topic 6) Iodine + Propanone + acid
Use Neuss handout or ILPAC book has even more detailed resource.
Explain principles of colorimetry
Explain calibration using test tube standard samples
Run through procedure
Explain reaction, gets lighter (transmittance/absorbance)
Set the program to record appropriate data
Amounts of Iodine and Water for Calibration of Colorimeter
Iodine
Water
5
5
1
9
0.5
9.5
0.5
19.5
0
10
Graphs tend to slope down like a hill for some reason. The reaction appears to go
FASTER as time progresses, instead of slower. This may be because of the colour
settling at the bottom of the cuvette, not stirring to ensure proper mixing, etc.
Calculations (instructions already given on the lab handout):
Plot absorbance vs. time (or transmittance vs. time) graphs for each trial, and find
initial tangents (rates). Use these tangents and your calibration curve to convert the
Absorbance/Transmittance into [I2] values. With [I2] rates, you can determine the
initial rate of the reaction for varying the concentrations of I2, HCl, and propanone.
From this, you can determine the order of the reaction wrt each of those components
Design Lab:
Investigate a factor that affects the rate of a chemical reaction,
but cannot be concentration, surface area, or temperature.
(Topic 9) Redox Reactions – Reactivity
In this lab, you will performing several simple experiments in test tubes. These
reactions will be simple displacement reactions that you have investigated before, but
never from a redox point of view. Read through the handout before you start
performing any of the experiments (some have to be performed early because it takes
time for the reaction to occur).
For each experiment:
- record observations
- write a balanced chemical equation
- indicate oxidation states of all elements in reactants and products
- indicate which elements were oxidized/reduced
Experiment 1: Reaction of Halides with Silver Nitrate (do this first)
1. Put about 3mL of the unknown halide solutions into 3 different test tubes
2. Add 1 pipette full of silver nitrate to the test tube#
3. Ensure the solutions are properly mixed, then observe to determine which
halide is Cl, Br, and I
Experiment 2: Metal Displacement Reactions
1. You will be performing these reactions in test tubes. Use no more than 3mL
of each solution, or a few grains of solid (when using solids).
a. CuSO4 + solid Mg
b. FeSO4 + solid Mg
c. CuSO4 + solid Fe
d. CuSO4 + solid Pb
e. MgSO4 + solid Cu
f. ZnSO4 + solid Cu
g. KI + solid Mg
h. KI + solid Fe
Experiment 3: Halogen Displacement Reactions
1. Obtain 3mL of halide solutions into each test tube.
2. Add 2mL of Br2 water to each solution, then add a few drops of hexane
3. Add 2mL of Cl2 water to each solution, then add a few drops of hexane
4. If you do not observe anything, wait for several minutes, then come back to it
Experiment 4: Reaction of Alkali metals with water (do Na and K inside the fume
hood)
1. Carefully cut a rice-sized piece of Li metal and react it in a large test tube 1/3
full of water. Do not use a piece that is too large
2. Repeat for Na and K metals.