precipitation reaction experiment report

Golden Rain Experiment

Lead Nitrate + Potassium Iodide

Lead nitrate reacts with potassium iodide to produce a beautiful precipitate, as we will show you. The reaction, known as the “Golden Rain” experiment , produces beautiful hexagonal crystals of lead iodide that resemble plates of gold , and makes a great chemistry demonstration.

The golden rain reaction takes advantage of the increased solubility of lead iodide in hot water. Stoichiometric amounts of lead nitrate and potassium iodide are combined, with enough water to dissolve all of the lead iodide precipitates at 80 degrees Celsius. When the solution cools, beautiful lead iodide crystals will fall out of solution.

Lead iodide golden rain experiment requirements

Lead (II) nitrate 1.65 grams (.005 moles) Potassium iodide 1.66 grams (.01 moles) Erlenmeyer flask 1000ml Hotplate-stirrer

Golden Rain Procedure – Tips & Tricks

Lead nitrate and potassium iodide are both solid, soluble ionic compounds. We will combine them for some amazing results.

• Dissolve each salt in 400ml of distilled water in separate beakers.
• Combine the liquids in the Erlenmeyer flask so you have 800ml in total. If you wish to use a 500 ml flask instead, simply cut the amounts of compounds and water in half. You will see a yellow precipitate of lead iodide fall out of solution.

Mastering chemistry challenge: How would you calculate the amounts needed yourself? Leave your answer in the comments.

PbI 2 will immediately precipitate out, as it is insoluble in cold water.

3. Heat the solution until all of the lead iodide dissolves, you may need to heat it above 80 degrees Celsius. Heating the solution causes the solubility to increase just enough to dissolve all of the lead iodide.

Lead iodide precipitate – how to best view it

4. Let it cool. This time, the PbI 2 precipitates out in a much more beautiful fashion . This is best viewed in a dark with bright sunlight shining onto the flask, for example through a garage window in the late afternoon. If the lead iodide settles too quickly, stir it with a long stirring rod or start magnetic stirring to keep the particles suspended – giving the “golden rain” effect.

The Golden Rain reaction

Here is the equation for this double-replacement reaction . Lead ii nitrate reacts with potassium iodide forming lead (II) iodide and potassium nitrate.

Pb(NO 3 ) 2 + 2KI -> PbI 2 + 2KNO 3 Net ionic equation : Pb +2 + 2I – -> PbI 2 (s)

Interesting fact: Lead is in the +2 oxidation state in this reaction. Lead (IV) iodide does not exist, because lead (IV) can oxidize iodide to iodine .

Lead / iodine complexes

Don’t use too much iodide, or this reaction will occur, forming the soluble colorless tetraiodoplumbate(II) complex ion.

PbI 2 + 2I – -> PbI 4 -2

Safety & Disposal

Lead nitrate is toxic, the lethal oral dose is approximately 8 grams for an 80kg human. Do not ingest any and avoid skin contact or breathing the dust.

The lead iodide should be filtered and stored in your compound collection. Lead salts should not be washed down the drain. The remaining lead in the solution can be precipitated out with sodium sulfide, as lead sulfide is extremely insoluble. PbS should be stored in a hazardous waste drawer until it can be disposed of properly. Sodium carbonate can be used if a sulfide compound is not available.

If you are doing a lab report, here is an example .

Lead Iodide / Golden Rain experiment video

Here’s the complete experiment video

We filmed this short clip on the golden rain reaction to show how beautiful the flakes of lead iodide look in the sun, when they precipitate in the cooled down solution. The video was taken in a dark garage, with sunlight coming through a window.

About Lead Iodide

Lead (II) iodide is a bright yellow solid, that is slightly soluble in hot water. It is stable in air. The formula for lead iodide is PbI 2 , and its molar mass is 461.01 grams/mole. The symbol for lead is Pb because its latin name is plumbum.

Lead iodide is quite a heavy molecule from a molar mass perspective, because both lead and iodine are heavier atoms. It has a hexagonal close-packed crystal structure, which is why it crystallizes in thin hexagonal-shaped plates. If you love math and crystals, read this.

Lead iodide is used in the manufacturing of solar cells, and also as a photon-detector for x-rays and gamma rays.

Related Articles

If you enjoyed this article about lead iodide and the golden rain experiment, check out making tin crystals , and read about standard reduction potentials and how to name ionic compounds

Extracting potassium metal from a banana

The element potassium

The Loaded Element Lead

Copper crystals made from zinc metal

Solubility Rules

Your browser is not supported

Sorry but it looks as if your browser is out of date. To get the best experience using our site we recommend that you upgrade or switch browsers.

Find a solution

• Skip to main content
• Skip to navigation

• Back to parent navigation item
• Primary teacher
• Secondary/FE teacher
• Early career or student teacher
• Higher education
• Curriculum support
• Literacy in science teaching
• Periodic table
• Interactive periodic table
• Climate change and sustainability
• Resources shop
• Collections
• Post-lockdown teaching support
• Remote teaching support
• Starters for ten
• Screen experiments
• Assessment for learning
• Microscale chemistry
• Faces of chemistry
• Classic chemistry experiments
• Nuffield practical collection
• Anecdotes for chemistry teachers
• On this day in chemistry
• Global experiments
• PhET interactive simulations
• Chemistry vignettes
• Context and problem based learning
• Journal of the month
• Chemistry and art
• Art analysis
• Pigments and colours
• Ancient art: today's technology
• Psychology and art theory
• Art and archaeology
• Artists as chemists
• The physics of restoration and conservation
• Ancient Egyptian art
• Ancient Greek art
• Ancient Roman art
• Classic chemistry demonstrations
• In search of solutions
• In search of more solutions
• Creative problem-solving in chemistry
• Solar spark
• Chemistry for non-specialists
• Health and safety in higher education
• Analytical chemistry introductions
• Exhibition chemistry
• Introductory maths for higher education
• Commercial skills for chemists
• Kitchen chemistry
• Journals how to guides
• Chemistry in health
• Chemistry in sport
• Chemistry in your cupboard
• Chocolate chemistry
• Adnoddau addysgu cemeg Cymraeg
• The chemistry of fireworks
• Festive chemistry
• Education in Chemistry
• Teach Chemistry
• On-demand online
• Live online
• Selected PD articles
• PD for primary teachers
• PD for secondary teachers
• What we offer
• Chartered Science Teacher (CSciTeach)
• Teacher mentoring
• UK Chemistry Olympiad
• Who can enter?
• How does it work?
• Resources and past papers
• Top of the Bench
• Schools' Analyst
• Regional support
• Education coordinators
• RSC Yusuf Hamied Inspirational Science Programme
• RSC Education News
• Supporting teacher training
• Interest groups

• More navigation items

Precipitation reactions of lead nitrate

Compare the colours of lead compounds formed by precipitation reactions to identify which would make good pigments in this microscale class practical

Many lead compounds are insoluble and some of them are brightly coloured. In this experiment, students observe the colour changes of lead nitrate solutions when different anions are added to identify which compounds would make good pigments.

They then compare what happens when deionised water and tap water are added to lead nitrate solution, giving explanations for what they observe.

The experiment should take approximately 20 minutes.

• Eye protection
• Student worksheet (available for download below)
• Clear plastic sheet (eg ohp sheet)

Solutions should be contained in plastic pipettes. See the accompanying  guidance on apparatus and techniques for microscale chemistry , which includes instructions for preparing a variety of solutions.

• Sodium hydroxide, 1 mol dm –3
• Lead nitrate, 0.5 mol dm –3
• Potassium iodide, 0.2 mol dm –3
• Sodium chloride, 0.5 mol dm –3
• Potassium bromide, 0.2 mol dm –3
• Sodium carbonate, 0.5 mol dm –3
• Sodium sulfate, 0.5 mol dm –3
• Potassium chromate, 0.2 mol dm –3

Health, safety and technical notes

• Read our standard health and safety guidance.
• Wear eye protection throughout (splash-resistant goggles to BS EN166 3).
• Sodium hydroxide solution, NaOH(aq), 1 mol dm –3  is CORROSIVE. See CLEAPSS Hazcard HC091a and CLEAPSS Recipe Book RB085.
• Lead nitrate, Pb(NO 3 ) 2 (aq), 0.5 mol dm –3  is a reproductive toxin and a specific target organ toxin. It also causes eye damage and is a probable carcinogen. See CLEAPSS Hazcard HC057a and CLEAPSS Recipe Book RB053.
• Potassium chromate, K 2 CrO 4 , 0.2 mol dm –3  is a carcinogen, mutagen and skin sensitiser. It is also toxic to aquatic life. Wear splash-proof eye-protection if transferring large amounts. Avoid skin contact. See CLEAPSS Hazcard HC078a and CLEAPSS Recipe Book RB069.
• Potassium bromide, KBr(aq), 0.2 mol dm –3 – see CLEAPSS Hazcard HC047b and CLEAPSS Recipe Book RB068.
• Sodium sulfate, Na 2 SO 3 (aq), 0.5 mol dm –3 – see CLEAPSS Hazcard HC098B and CLEAPSS Recipe Book RB107.
• Sodium carbonate, 0.5 mol dm –3 – see CLEAPSS Hazcard HC095A and CLEAPSS Recipe Book RB080.
• Potassium iodide, KI(aq), 0.2 mol dm –3 – see CLEAPSS Hazcard HC047b and CLEAPSS Recipe Book RB072.

Part 1: adding different anions to lead nitrate solution

• Cover the worksheet with a clear plastic sheet.
• Put one drop of lead nitrate solution in each box of table 1.
• Add one drop of each of the solutions containing the anions indicated to the appropriate box.

Part 2: adding deionised water and tap water to lead nitrate solution

• With the worksheet still covered, put one drop of lead nitrate solution into each box of table 2.
• Add one drop of deionised water and one drop of tap water to the appropriate boxes.

Questions for students

• Which of the lead compounds observed appear to be good pigments?
• What is the main disadvantage of using these compounds as pigments?

What explanations can you give for your observations?

Teaching notes and expected observations

The addition of solutions of each of the anions produces precipitates, which indicates that in general lead compounds are insoluble. The iodide is an intense yellow colour, the chromate(VI) is also yellow and both could be used as pigments except for the fact that lead compounds are toxic.

The fact that lead forms insoluble compounds is used as a basis for indicating the presence of anions in water. The addition of deionised water to lead nitrate gives no cloudiness. However, with tap water a cloudiness gradually develops if the water is from a hard water area since carbonates, sulphates or hydrogen carbonates may be present. If you live in a soft water area there will probably be no cloudiness. (One solution is to simulate hard water conditions.)

Precipitation reactions of lead nitrate - student sheet

Precipitation reactions of lead nitrate - teacher notes, additional information.

This resource is part of our  Microscale chemistry  collection, which brings together smaller-scale experiments to engage your students and explore key chemical ideas. The resources originally appeared in the book  Microscale chemistry: experiments in miniature , published by the Royal Society of Chemistry in 1998.

© Royal Society of Chemistry

Health and safety checked, 2018

• 14-16 years
• 16-18 years
• Practical experiments
• Properties of matter
• Reactions and synthesis

Specification

• 3.19 Recall the general rules which describe the solubility of common types of substances in water: all common sodium, potassium and ammonium salts are soluble; all nitrates are soluble; common chlorides are soluble except those of silver and lead…
• 3.20 Predict, using solubility rules, whether or not a precipitate will be formed when named solutions are mixed together, naming the precipitate if any
• (g) nature, physical properties and acid-base properties of CO₂ and PbO
• (i) reactions of Pb²⁺(aq) with aqueous NaOH, Cl⁻ and I⁻

Related articles

Non-burning paper: investigate the fire triangle and conditions for combustion

2024-06-10T05:00:00Z By Declan Fleming

Use this reworking of the classic non-burning £5 note demonstration to explore combustion with learners aged 11–16 years

Everything you need to introduce alkenes

2024-06-04T08:22:00Z By Dan Beech

Help your 14–16 learners to master the fundamentals of the reactions of alkenes with these ideas and activities

Particle diagrams | Structure strip | 14–16

By Kristy Turner

Support learners to describe and evaluate the particle model for solids, liquids and gases with this writing activity

1 Reader's comment

Only registered users can comment on this article., more experiments.

‘Gold’ coins on a microscale | 14–16 years

By Dorothy Warren and Sandrine Bouchelkia

Practical experiment where learners produce ‘gold’ coins by electroplating a copper coin with zinc, includes follow-up worksheet

Practical potions microscale | 11–14 years

By Kirsty Patterson

Observe chemical changes in this microscale experiment with a spooky twist.

Antibacterial properties of the halogens | 14–18 years

Use this practical to investigate how solutions of the halogens inhibit the growth of bacteria and which is most effective

• Contributors
• Email alerts

Site powered by Webvision Cloud

Skip to Main Content

• My Assessments
• My Curriculum Maps
• Communities
• Workshop Evaluation

Share Suggestion

Precipitation reaction lab, grade levels, course, subject.

• Printer Friendly Version

Predict how combinations of substances can result in physical and/or chemical changes.

Interpret and apply the laws of conservation of mass, constant composition (definite proportions), and multiple proportions.

Balance chemical equations by applying the laws of conservation of mass.

Classify chemical reactions as synthesis (combination), decomposition, single displacement (replacement), double displacement, and combustion.

Use stoichiometry to predict quantitative relationships in a chemical reaction.

1. Students will determine which combinations of ionic solutions form precipitates.

2. Students will identify the precipitate in each reaction.

3. Students will write net ionic equations for each reaction.

4. Students will write patterns or trends for certain cations and anions.

1. Construct a data table to record observations. Please list the chemicals in alphabetical order as above.

2. Add 1-2 drops of each pair of chemicals to a depression well on your spot plate.

3. If a precipitate forms, record that data including color.

4. If no reaction occurs, record N.R. on your data table.

5. Repeat procedures #1-4 for all chemical pairs listed on your data table.

Description

The majority of ionic solids are soluble in water. Those that are not, form solid products called precipitates when two aqueous ionic solutions are mixed.

Recall that ionic compounds are made of positive and negative ions held together by the attractive, electrostatic forces that occur between oppositely charged particles. Soluble ionic compounds break apart completely into their respective ions when put in water. Example: NaCl (s) when put into water yields Na + (aq) and Cl - (aq) , and Ag(NO) 3 (s) also dissociates in water to form these respective ions - Ag + (aq) and N0 3 - (aq) . It turns out that when these two solutions of sodium chloride and silver nitrate are mixed a solid falls out (precipitation occurs). The new mixture still contains Na + (aq) and N0 3 - (aq) , however, the newly formed precipitate is AgCl (s) . The chemist describes this process first as a complete ionic equation .

Na + (aq) + Cl - (aq) + Ag + (aq) + NO 3 - (aq) ---> AgCl (s) + Na + (aq) + NO 3 - (aq)

Notice that the sodium and nitrate ions appear the same on both sides of the equation, that means they did not change; therefore, they are called spectator ions. Chemists like to write a more useful equation that describes only the changes that took place. They write a net ionic equation , which eliminates spectator ions.

NET: Ag + (aq) + Cl - (aq) -----> AgCl (s)

In this experiment, you will mix 13 different ionic solutions in all possible combinations to determine which combinations result in precipitate formation. Based on your results, you will write complete and net ionic equations for each reaction that has taken place. Use your solubility chart to determine what precipitates.

Dropper bottles with the following 0.20 M solutions.

1. barium chloride 7. potassium carbonate

2. cobalt chloride 8. potassium hydroxide

3. copper (II) sulfate 9. potassium iodide

4. iron (III) nitrate 10. silver nitrate

5. lead (II) nitrate 11. sodium carbonate

6. nickel (II) sulfate 12. sodium sulfide

Disposal and Safety

1. Rinse your spot plates into the waste container provided.

2. Use Q-tips and soap to clean out each well, rinse with tap water.

3. Use distilled water as your final rinse.

4. Clean up your lab area and reagent table. Wash your hands thoroughly.

Data Analysis

1. For each combination of solutions that GAVE A PRECIPITATE, write the correct net Ionic equation. Use a solubility chart to determine the identity of the precipitate.

2. Write these net ionic reactions in alphabetical order as they appear above. For example, the first reaction should be those containing barium chloride.

3. Carefully examine your net ionic reactions and data table, what patterns do you notice for certain precipitate reactions? What chemicals do not precipitate?

Date Published

Insert template.

Investigating Precipitation Reactions: Lab Report & Results

• school Campus Bookshelves
• menu_book Bookshelves
• perm_media Learning Objects
• login Login
• how_to_reg Request Instructor Account
• hub Instructor Commons

Margin Size

• Download Page (PDF)
• Download Full Book (PDF)
• Periodic Table
• Physics Constants
• Scientific Calculator
• Reference & Cite
• Tools expand_more
• Readability

selected template will load here

This action is not available.

7: Gravimetric Analysis (Experiment)

• Last updated
• Save as PDF
• Page ID 95886

• Santa Monica College

\( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

\( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)

\( \newcommand{\id}{\mathrm{id}}\) \( \newcommand{\Span}{\mathrm{span}}\)

( \newcommand{\kernel}{\mathrm{null}\,}\) \( \newcommand{\range}{\mathrm{range}\,}\)

\( \newcommand{\RealPart}{\mathrm{Re}}\) \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)

\( \newcommand{\Argument}{\mathrm{Arg}}\) \( \newcommand{\norm}[1]{\| #1 \|}\)

\( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)

\( \newcommand{\Span}{\mathrm{span}}\)

\( \newcommand{\id}{\mathrm{id}}\)

\( \newcommand{\kernel}{\mathrm{null}\,}\)

\( \newcommand{\range}{\mathrm{range}\,}\)

\( \newcommand{\RealPart}{\mathrm{Re}}\)

\( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)

\( \newcommand{\Argument}{\mathrm{Arg}}\)

\( \newcommand{\norm}[1]{\| #1 \|}\)

\( \newcommand{\Span}{\mathrm{span}}\) \( \newcommand{\AA}{\unicode[.8,0]{x212B}}\)

\( \newcommand{\vectorA}[1]{\vec{#1}}      % arrow\)

\( \newcommand{\vectorAt}[1]{\vec{\text{#1}}}      % arrow\)

\( \newcommand{\vectorB}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

\( \newcommand{\vectorC}[1]{\textbf{#1}} \)

\( \newcommand{\vectorD}[1]{\overrightarrow{#1}} \)

\( \newcommand{\vectorDt}[1]{\overrightarrow{\text{#1}}} \)

\( \newcommand{\vectE}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash{\mathbf {#1}}}} \)

• To experimentally analyze an unknown sulfate salt via a precipitation reaction, using the techniques associated with Gravimetric Analysis to collect and weigh the precipitate, and
• To calculate the percentage by mass of \(\ce{SO_4^{-2}}\) in the unknown sulfate salt via a stoichiometric analysis of the collected precipitate, and then use this percentage to identify the metal “M” present in the sulfate salt.

Gravimetric analysis is a quantitative method for accurately determining the amount of a substance by selective precipitation of the substance from an aqueous solution. The precipitate is separated from the remaining aqueous solution by filtration and is then weighed. Assuming that the chemical formula for the precipitate is known and that the precipitation reaction goes all the way to completion, then the mass of the substance in the original sample can be determined.

In this experiment, the percentage by mass of sulfate in an unknown sulfate salt will be determined by gravimetric analysis. First, a pre-weighed sample of the unknown sulfate salt will be dissolved in water. Next, an excess of aqueous barium chloride is added to the aqueous solution of the unknown salt. This will result in the precipitation of all the sulfate ions as barium sulfate:

\[ \text{Metal sulfate (aq) + Barium chloride (aq)} \ce{->} \text{Barium sulfate (s) + Metal chloride (aq)}\]

The barium sulfate precipitate is collected by filtration, dried and weighed. Since barium chloride is added in excess, and since the precipitation reaction goes to completion, we can assume that all of the sulfate is transferred from the original unknown sample to the precipitate. The mass of sulfate in the collected \(\ce{BaSO4}\) precipitate can be calculated via its percent composition. This also yields the mass of sulfate in the original unknown since:

\[ \text{mass of sulfate in the precipitate = mass of sulfate in the unknown sample}\]

Finally, using the mass of sulfate along with the initial mass of unknown used, the percentage by mass of sulfate in the original sample may now be calculated.

In order to obtain the best results, the collected \(\ce{BaSO4}\) crystals should be as large as possible. This considerably aids the filtration process (larger crystals are less likely to be pass through the filter paper), and it also minimizes the amount of impurities adsorbed onto the crystals (smaller surface area). In general, larger crystals are obtained when the rate of precipitation is as low as possible. The rate of precipitation is minimized by slowly adding the \(\ce{BaCl2}\) solution to the aqueous mixture containing the unknown salt while continuously stirring the mixture. The rate of precipitation can be decreased even further by slightly increasing the solubility of the \(\ce{BaSO4}\). This may be achieved by lowering the pH with 6 M \(\ce{HCl}\) and by increasing the temperature. The resulting decrease in the yield of the \(\ce{BaSO4}\) is insignificant.

Unknown solid sulfate salt, 6 M \(\ce{HCl}\) solution, and 0.1 M \(\ce{BaCl2}\) solution

250-mL beaker, analytical balance, stirring rod, 100-mL graduated cylinder, stand with ring clamp, wire screen, Bunsen burner, wash bottle with distilled water, crucible and lid, crucible tongs, ash-less filter paper, large funnel, 500-mL Erlenmeyer flask, clay triangle

Be very careful when handling 6 M \(\ce{HCl}\) ( aq ). If this acid comes in contact with your skin or eyes you should immediately rinse the affected area with water for several minutes. Also remember that items heated in the Bunsen burner are very hot (especially the crucible), and to allow ample time for them to cool before touching.

Instructions

• Weigh a clean, dry 250-mL beaker to the nearest 0.001 g using the analytical balance, and record this mass on your lab report. Next, add 0.30 – 0.35 grams of your unknown sample to the beaker. Record the combined mass of the beaker plus sample on your lab report.
• Add 50 mL of distilled water, followed by 20 drops of 6 M \(\ce{HCl}\) ( aq ), to the sample in the beaker. Stir the contents of the beaker until the sample has entirely dissolved. Leave the stirring rod in the beaker.
• Obtain a stand with a ring clamp from the back of the lab. Place your wire screen on the ring and the beaker containing your dissolved sample on the wire screen. Use the Bunsen burner to heat the solution until it is nearly (but not quite) boiling. Turn the Bunsen burner off before the solution boils.
• While heating the solution, measure out 25 mL of 0.1 M \(\ce{BaCl2}\) ( aq ) using a 100-mL graduated cylinder. The graduated cylinder used should be clean (rinse with distilled water) but does not need to be dry.
• Slowly add small portions of \(\ce{BaCl2}\) ( aq ) to the beaker containing the hot solution. You should observe the formation of a white precipitate of \(\ce{BaSO4}\) ( s ). Stir the contents of the beaker as you add the \(\ce{BaCl2}\) solution. The addition of the \(\ce{BaCl2}\) must be performed very slowly – this step should take you at least 3 minutes to complete! When finished, rinse any precipitate that remains on the stirring rod into the solution with a small amount of distilled water, and then allow the precipitate to settle in the beaker for about 20 minutes.
• While the precipitate settles, prepare your crucible by heating it in the hottest part of the Bunsen burner flame for about 2 minutes (use the crucible tongs, cradling the crucible as demonstrated by the instructor – never pick up a crucible by pinching the walls ). Repeat with the lid. Place the hot crucible and lid on the metal base of the stand to cool. Once they have cooled to room temperature weigh the crucible without the lid using the analytical balance, and record this mass on your lab report. You do not need to weigh the lid.
• Obtain a piece of ash-less filter paper from your instructor and fold it into quarters. Open the folded paper into a cone and place it into your large funnel. Wet the filter paper with a small amount of distilled water so that it adheres to the funnel. Sit the funnel in the mouth of a 500-mL Erlenmeyer flask, which will be used to collect the filtrate.
• After 20 minutes has passed, slowly pour the mixture containing the \(\ce{BaSO4}\) precipitate down your stirring rod into the funnel. Be careful that the level of liquid in the funnel is never more than three-fourths of the way to the top of the filter paper. When the transfer is complete use your wash bottle (filled with distilled water) to rinse the residual precipitate from the beaker and the stirring rod into the funnel.
• After all the liquid has drained from the funnel, wearing gloves very carefully press the top edges of the filter paper together, and gently fold the filter paper into a compact package that will fit into the crucible. It is important that you do not use too much force in order to avoid tearing the filter paper. Place the folded filter paper into the crucible.
• Take your stand, ring clamp and Bunsen burner over to a fume hood. Place your clay triangle on the ring and the crucible in the clay triangle for support. Gently heat the crucible without the lid to remove the water. Once the paper appears to be dry (after several minutes), heat the crucible more vigorously so that the filter paper begins to char (turning from white, to brown, to black) – but not so vigorously that the filter paper bursts into flame. If the filter paper bursts into flame you should cover it with the crucible lid to put out the flame, then reduce the amount of heat and remove the lid. Continue to heat moderately with the lid off until all of the filter paper has turned black.
• Once all the filter paper has turned black, vigorously heat the crucible without the lid in the hottest part of the Bunsen burner flame so that the bottom of the crucible is red hot . The center of the flame should now be directly on the crucible. The charred filter paper (carbon) will gradually combust and be converted into \(\ce{CO2}\) gas. When the filter paper is entirely combusted only the white \(\ce{BaSO4}\) should remain in the crucible. Continue to heat the crucible vigorously until no charred filter paper remains. This should take 5 - 10 minutes.
• Allow the crucible to cool to room temperature (this takes at least 5 minutes). Weigh the crucible without the lid and its contents on the analytical balance. Record this mass.
• Place the crucible and its contents back in the clay triangle and heat vigorously with the lid off for an additional 5 minutes. Then allow it to cool again and reweigh the crucible and its contents without the lid . If the mass is within 0.005 grams of that in step 12, then record this mass on your lab report. If the mass has decreased by more that 0.005 grams, then either the \(\ce{BaSO4}\) is still wet or not all of the filter paper has combusted and you should repeat this step until you achieve a consistent mass, making sure the crucible is in the hottest part of the flame.
• Discard the \(\ce{BaSO4}\) in the proper waste container, then clean up as directed by your instructor.

Pre-laboratory Assignment: Gravimetric Analysis

Suppose that 0.323 g of an unknown sulfate salt is dissolved in 50 mL of water. The solution is acidified with 6 M \(\ce{HCl}\), heated, and an excess of aqueous \(\ce{BaCl2}\) is slowly added to the mixture resulting in the formation of a white precipitate.

• Assuming that 0.433 g of precipitate is recovered calculate the percent by mass of \(\ce{SO4^{2-}}\) in the unknown salt.
• If it is assumed that the salt is an alkali sulfate determine the identity of the alkali cation.

Lab Report: Gravimetric Analysis of an Unknown Sulfate

Experimental Data

• Unknown Sulfate ID Code:

Calculations and Conclusions

• Calculate the mass of sulfate in the unknown sample. Clearly show each step of your calculation.
• Calculate the percentage by mass of sulfate in the unknown sample. Show your work.
• Calculate the percentage by mass of the metal in the unknown sample. Show your work.
• The cation in your unknown metal sulfate is one of the following:
• \(\ce{Al^{3+}}\)
• \(\ce{Na^+}\)
• \(\ce{Ni^{2+}}\)
• \(\ce{K^+}\)
• \(\ce{NH^{4+}}\)
• \(\ce{Cd^{2+}}\)

Use this information along with your experimental results to determine which cation it is.

• Identity of Cation:

Show all of your work with clear, logical steps below. Clearly explain how your calculations here along with your experimental results for #2 and/or #3 allowed you to identify the cation in the metal sulfate.

• Suppose an unknown metal sulfate is found to be 72.07% \(\ce{SO4^{2-}}\). Assuming the charge on the metal cation is +3, determine the identity of the cation.
• The unknown metal sulfates are hygroscopic and will absorb water from air. The unknowns must thus be kept in desiccators to remove any absorbed water. How would your results be affected if your unknown sample was not desiccated? Would this error cause your calculation of the mass percent of sulfate in the unknown to be too high or too low? Explain.
• In this experiment you used an excess of the \(\ce{BaCl2}\) solution. How would your results be affected if you did not use an excess of the \(\ce{BaCl2}\) solution? Would this error cause your calculation of the mass percent of sulfate in the unknown to be too high or too low? Explain.
• In the last step of the procedure, you vigorously heated the \(\ce{BaSO4}\) precipitate wrapped in filter paper in a crucible. How would your results be affected if tiny pieces of the filter paper still remained mixed in with the \(\ce{BaSO4}\) after heating? Would this error cause your calculation of the mass percent of sulfate in the unknown to be too high or too low? Explain.

Verify originality of an essay

Get ideas for your paper

Find top study documents

Finding The Limiting Reagent In A Chemical Reaction: Lab Report

Published 20 Jun 2024

1. Introduction

The purpose of this investigation is to gain quantitative information of the reaction that occurs when Sodium iodide and Lead (II) nitrate are combined.

Stoichiometry is the quantitative relation between the number of moles (and therefore mass) of various products and reactants in a chemical reaction (Washington University in St. Louis, 2005.). Chemists use stoichiometry as they are responsible for designing a chemical reaction and analysing the products obtained from it. Chemists must determine the amount of each reactant that is required and the amount of each product that will be produced (Dr. Bailey, n.d.).

The limiting reagent, in a chemical reaction, is the reactant that determines how much of the products are made. The Excess Reagent, in a chemical reaction, is the left-over reactant once the limiting reagent is completely consumed (Khan Academy, 2019). Once all the limiting reagent has been used the reaction cannot continue.

Theoretical yield is the quantity of a product produced from the entire consumption of a limiting reactant in a chemical reaction. It is the amount of product resulting from a perfect (theoretical) chemical reaction (Helmenstine, 2019). Theoretical yield can calculated by first balancing the equation, identifying the limiting reactant, understand the mole ratio and the moles of the products then multiplying the moles in conjunction with the ratio and finally multiply those moles by the molecular weight of each product to get the mass in grams. The actual yield is the amount of the substance that is actually produced in a reaction. It is very unlikely that the actual yield and theoretical yield will be the same, this is due to possible transferring error and or measuring error among others.

In this experiment, the fixed mass of Lead (II) nitrate is being reacted with incremental masses of Sodium iodide. The yellow precipitate (Lead iodide) is used as a measurement for the limiting reagent analysis in the influence of the amount of product that is formed. The prediction of a maximum of 2.31g of Lead iodide can be calculated through the theoretical yield and the balanced chemical equation (refer to equation 1 & 2) shows that the mole ratio of 1:2 is apparent in this equation.

Pb(NO3)2 + 2NaI PbI2 + 2NaNO3 Eq. 2

The aim of this experiment was to conduct a chemical reaction, which created Lead iodide precipitate to be gathered and analysed, and to find the limiting reagent within the reaction.

Read also: Saying, " do my lab report "? Our professional team will handle it, delivering accurate and timely results.

3. Hypothesis

It can be hypothesised that the limiting reagent of this experiment is Sodium iodide and having a constant of 1.66 grams of Lead (II) nitrate means that the maximum increment of 1.50 grams of Sodium iodide can be reacted to form a theoretical yield of 2.31 grams of Lead iodide. This is because of the Sodium iodide to Lead nitrate and Sodium iodide to Lead iodide mole ratio in the balanced equation is 2 is to 1. Due to this mole ratio it is proven that the maximum amount of reactant and the maximum amount of product with the constant of 1.66 grams of Lead (II) nitrate is 1.50 grams of NaI and 2.31 grams of Lead iodide. (See appendix for calculation).

4. Materials

• 100ml beaker (x 20)
• 100ml conical flask (x 10)
• Filter funnel (x 1)
• Stirring rod (x 1)
• Wash bottle – deionised water (x 1)
• Ethanol – 10ml
• 2-place balance (x 1)
• Oven at 110◦C (Fahrenheit)/ 35◦C (Degrees)
• Disposable gloves
• Safety glasses
• Lead nitrate – 20 grams
• Sodium iodide – 30 grams

5. Variables

The following table shows the various variables involved in this experiment. These are: Independent Variable, Dependent Variable and Controlled Variables

Description/Reasoning

Independent variable

Sodium iodine (NaI) mass (g)

Used NaI in 10 separate increments of 0.15g starting at 0.75g

Dependent variable

Lead iodine (PbI2) mass (g)

The mass of PbI2 depended on the change of NaI and each reaction was measured.

Controlled variables

Lead (II) nitrate {Pb(NO3)2} mass (g)

1.66g of Pb(NO3)2 was used in each experiment

Water (H2O)

30ml used in each experiment

Balancing equipment (2-place)

Each experiment used the same balancer

Precipitation waiting time

All 3 - 5 minutes

Drying time

All kept overnight

Temperature

all in a 100◦C (Fahrenheit)/35◦C (degrees) oven overnight

Conditions of the experiment

Experimental process on the same day as each other

6. Safety Measures

This table shows the shows the safety measures considered prior and during the experiment.

Safety Precaution

• Safety Glasses

Skin Damage

• Hard covered shoes
• Do not look directly over reactions
• do not smell products (wafting technique)
• be careful with observations

Poisoning/Consumption

• Handle with extreme caution,
• dispose immediately of waste once used, wear gloves,
• do not consume

Equipment/Glass ware

Glass Breakage

• Handle glassware with care,
• place away from edge when not in use

Injury to skin/cuts

• handle any broken glass with care
• handle all things with care
• The Equipment and chemicals were collected and all safety measures adhered to.
• The filter paper was folded into quarters and then it was weighed. The weight of the filter paper was then written down on the board in a table with the corresponding experiment number.
• Both the Lead (II) nitrate and the NaI was measured out on weigh boats to the closest possible reading.
• The mass of the Lead (II) nitrate was written down on both sides of the filter paper at the open end of the fold.
• 30ml of deionised water was added to each of the 100ml beakers.
• Each measured reactant was transferred into separate beakers. Then the weigh boats were rinsed with deionised water and added to ensure no transferring error occurred.
• Each solution was then stirred with the stirring rod.
• The Lead (II) nitrate solution was then carefully added to the Sodium iodine mixture. Then the beaker was rinsed out with the deionised water 3 times to ensure no transferring error occurred.
• The new Lead (II) nitrate and Sodium iodine mixture was stirred and then left for 3 - 5 minutes.
• The filter paper was then placed into a filter funnel which was sitting in a 100ml conical flask. Then deionised water was filtered through to ensure the paper stayed in place.
• The precipitated solution was then transferred through the filter paper creating filtration and yellow residue.
• To ensure a complete transfer the beaker was rinsed multiple times with deionised water as well as ethanol.
• After the filtration process finalised the filtration paper was carefully removed from the funnel and placed into a beaker by the instructor.
• The instructor then placed all of the beakers into the oven which was 100◦C (Fahrenheit)/35◦C (degrees) and it was left overnight.
• The next morning the instructor then removed the paper and allowed to cool then began reweighing the filter paper and precipitate. The difference of the filter paper was then found through deducting the original mass from the filter paper with the residue’s mass.

This table shows the moles and grams of Lead (II) nitrate and Sodium iodine used in the experiment. It also shows the theoretical, actual and percentage yield of lead iodine produced.

8. Refence List

• Washington University in St. Louis, 2005. http://www.chemistry.wustl.edu/~coursedev/Online%20tutorials/Stoichiometry.htm Dr. Kristy M. Bailey no date Professor of Chemistry
• Oklahoma City Community College http://www.occc.edu/kmbailey/Chem1115Tutorials/Stoichiometry_Map.htm
• Khan Academy 2019 https://www.khanacademy.org/science/chemistry/chemical-reactions-stoichiome/limiting-reagent-stoichiometry/a/limiting-reagents-and-percent-yield
• Anne Marie Helmenstine, Ph.D. updated July 14, 2019, https://www.thoughtco.com/theoretical-yield-definition-602125

Was this helpful?

Thanks for your feedback, related blog posts, the great controversy about the homework.

Imagine this. You are coming home from a long school day just wanting to relax, maybe take a long nap because you are tired from staying up late do...

Arguments Regarding 'No Homework Policy'

No homework policy is about banning all the teachers from giving homework or assignments. There are different reactions and opinions there are some...

Homeostasis Lab Report: How The Time Of The Exercise Affects Heart Rate During Exercise

Introduction Homeostasis is the tendency toward a relatively stable equilibrium between interdependent elements. Homeostasis refers to stability...

Join our 150K of happy users

• Get original papers written according to your instructions
• Save time for what matters most