LAB [25 pts]
Stoichiometry &
Precipitation of PbCrO4
NAME
Lab Partner
Period
Date
Introduction: When doing a reaction, it is important to know how much of each reactant is needed to form a
particular amount of product. In this lab, you will be doing the following precipitation reaction:
K2CrO4 (aq)
+
Pb(NO3)2 (aq)
→
PbCrO4 (s)
+
2 KNO3 (aq)
potassium chromate
lead nitrate
lead chromate
potassium nitrate
(clear yellow)
(clear colorless)
(yellow solid)
(colorless solution- white solid)
You will compare how much product you actually isolated with how much you should have gotten (according
to mass-mass calculations). In addition, this lab will introduce you to some basic laboratory techniques such as
filtration, washing and drying.
Purpose: The purpose of the lab is to successfully carry out the above precipitation reaction, successfully isolate
the products and obtain a good yield of both products.
Prelab question: [1 pt]
If 1.00 g of K2CrO4 is reacted with Pb(NO3)2, how many grams of Pb(NO3)2 will be needed to completely
react with all of the K2CrO4? Show calculations here. (Mass → Mol → Mol → Mass)
Procedure:
1) Mass out close to 1.00 g of the yellow K2CrO4 (weighing boat). Carefully transfer the yellow K2CrO4 (s) into
a clean 250 ml beaker.
2) Mass out a clean, dry 250 mL Erlenmeyer flask. Record the exact value.
3) Mass out the amount of Pb(NO3)2 needed to react with all of the K2CrO4. (See your calculation above and then
increase that amount by 0.05g-0.09 g.) Put the white Pb(NO3)2 (s) in flask. Record mass.
4) Completely dissolve each of the solids by adding about 25 ml of distilled water to each solid. Carefully swirl
the solutions. Make sure that each solid has dissolved.
5) Slowly pour the Pb(NO3)2 solution (in flask) into the K2CrO4 solution in the beaker. Rinse the last traces of
Pb(NO3)2 from the flask into the beaker with a little distilled water (use water bottle). Carefully swirl the
resulting mixture (in beaker) for a good three minutes to make sure substances have reacted fully.
6) Obtain a piece a filter paper and write your names on it with pencil. Mass it out and record in table.
7) Set up ring stand with ring and funnel. Fold the filter paper, fit the filter paper into the funnel and moisten
the paper with some distilled water.
8) Put the flask under the funnel. Carefully pour the mixture from the beaker into the filter paper-- Don’t let it
overflow!!! (Hint: Swirl and pour quickly-- so that you get the maximum amount of solid out. )
9) Try to remove as much of the yellow solid from the beaker as possible by spraying the beaker with some
distilled water and pouring into filter. (It helps if you tilt your beaker above the filter as you spray.)
10) Wash the solid in filter paper by spraying with water bottle (Don’t overdo it-- 5 seconds is long enough.)
11) When the filtering is complete, carefully remove the filter paper from the funnel. Carefully unfold it and lay
it flat to dry overnight. (You will mass it out dry, tomorrow.)
12) Remove the flask and boil the filtrate (the solution in the flask) on a hot plate. It may take as long as 20 to
30 minutes to boil away all of the water, so take turns keeping watch. Remove flask from the hot plate
immediately when all the liquid has boiled away.
Day 2:
13) Mass out the dry filter paper with yellow residue (PbCrO4). Record mass.
14) Mass out your flask with the white residue (KNO3) in it. Record mass.
DATA: Record all data here. Record masses to nearest 0.01 g. Show ALL UNITS!!!
Reactants:
Products:
Mass of K2CrO4 solid (yellow)
Mass of dry filter paper & PbCrO4 (yellow)
Mass of dry PbCrO4 (yellow)
Mass of flask
Mass of Pb(NO3)2 solid (white)
Mass of flask and KNO3 (white)
Mass of filter paper
Mass of KNO3 (white)
Color of filtrate
(The filtrate is the solution that went through filter.)
Calculations: [6 pts] You only need to show work for calculations in # e-h.
a) Fill in this chart by copying down the experimental masses for each substance. Add up the molar masses for
each. Then convert each mass to moles (use their molar masses!!!!)
You should keep 3 SIG FIGS!
Experimental Mass (g)
Molar Mass (g/mole)
# of moles
Part (c)
K2CrO4 (yellow)
Pb(NO3)2 (white)
PbCrO4 (yellow)
KNO3 (white)
b) Compare the sum of the masses of the reactants to the sum of the masses of the products.
Total mass of reactants =
Mass lost/gained during reaction =
Total mass of products =
Did you lose or gain mass?
c) Determine if the coefficients of the balanced equation basically agree with your experimental mole values. To
do so, look at your mole values in your table above. Divide each one by the K2CrO4 mole value. Write in your answers in
column labeled, “Part (c)” in the chart above (use 3 sig figs).
Based on the balanced equation, what should the mole ratios be? 1 :
→
:
Do your experimental mole values basically agree with this? _______ If not, which are off a lot?_______
d) Based on the mole ratios you determined in (c), which reactant was the limiting reactant? ____________
Explain your reasoning: _______________________________________________________
e) Calculate the theoretical yield of PbCrO4 . (Assume that all of your limiting reactant reacts. Thus, use the grams of
your limiting reactant to determine how many grams of PbCrO4 should be produced.)
f) Calculate your percent yield of PbCrO4.
% yield =
experimental yield(g)
theoretical yield(g)
x100 =
g) Calculate the theoretical yield of KNO3. (Assume that all of your limiting reactant reacts. Thus, use the grams of
your limiting reactant to determine how many grams of KNO3 should be produced. )
h) Calculate your percent yield of KNO3.
The rest of the lab must be typed. Single spaced is fine.
Conclusions: [11 pts] You must answer in complete sentences.
1. [2 pts]As you know, mass should be conserved in a chemical reaction. However, sometimes there are
experimental errors that cause some gain or loss of mass. Did you lose or gain mass overall? Give two
possible experimental errors that could account for this overall change in mass.
Note: An experimental error cannot be an error due to measuring or calculating wrong. Also, the error cannot be
due to inaccurate or imprecise measuring equipment.
2. [1 pts] Based on your experimental mole ratios (See Calculations-part c), determine which reactant should
have been the excess reactant and explain your reasoning. What was the color of your filtrate (the solution
that went through your filter and into the flask)? Based on your color of the filtrate, which reactant was
actually in excess?
3. [4 pts] State the % yield of your PbCrO4. Was your yield too high or too low? No matter how good your
yield, discuss two experimental errors which could account for your high or low yield of PbCrO4.
4. [4 pts] State the % yield of your KNO3. Was your yield too high or too low? No matter how good your
yield, discuss two experimental errors which could account for your high or low yield of KNO3.
Post Lab Questions: [9 pts] (Please explain thoroughly and answer in COMPLETE sentences.)
1. [2 pts] In this lab, you were able to separate the PbCrO4( yellow product) from the KNO3 (white product).
Describe the basic process of isolating the products and explain how their solubility differences allowed for
isolation. (How was PbCrO4 collected? How was KNO3 collected? Relate each to solubility.)
2. [2 pts] In this experiment, you tried to add the exact mole ratios needed to allow both reactants to completely
react.
a) How would the purity of your two products have been affected if you had used excess K2CrO4? Which
product would be contaminated-- PbCrO4 or KNO3? Why would this product be contaminated (solubility!!)? Would you
be able to see this contamination (colors?)?
b) How would the purity of your two products have been affected if you had used excess Pb(NO3)2? Which
product would be contaminated-- PbCrO4 or KNO3? Why would this product be contaminated (solubility!!)? Would you
be able to see this contamination (colors?)?
3. This experiment should have convinced you that mass must be conserved in a chemical reaction, but moles
do not have to be conserved. (In other words, the mass of the reactants must equal the mass of the products. However, the
total moles of reactants do NOT have to equal the total moles of the products).
a) [1 pt] Explain why mass must be conserved in a chemical reaction. (Your explanation must include what
happens to atoms in chemical reactions.)
b) [2 pts] Explain why it is not necessary for moles to be conserved in a chemical reaction.
HINT: For the reaction in this lab, if you started with 2 moles, you would end up with 3 moles. Explain what this
means without using the term mole. Now, without mentioning mass or molar mass, explain why moles don’t have to be
conserved in a chemical reaction.
• To help you understand what is going on, you may want to explain how it can be possible to start with one mole and
form 2 moles.