HONORS CHEMISTRY
HARVARD-WESTLAKE
UNIT 3
Stoichiometry
Got mole problems? Call Avogadro at 602-1023.
UNIT 3
1
UNIT 3
2
Stoichiometry Problems
1. Iron metal does not occur in native form to any appreciable extent. Most of it is found in the
form of various oxides, iron (III) oxide being fairly common as an ore. This oxide is heated in
an atmosphere of carbon monoxide to produce iron metal. Carbon dioxide is also a product of
this process. How many grams of CO are needed to produce 1.0 kg of Fe?
2. One way to remove excess CO2 from recirculated air in a spacecraft is to react it with LiOH
to form lithium carbonate and water. How many grams of LiOH are needed for a 6.0 day 3person mission if each person exhales about 1.0 kg of CO2 each day?
UNIT 3
3
3. When nitroglycerin explodes it produces four gases:
C3H5(NO3)3  CO2 + N2 + O2 + H2O [not balanced]
How many grams of water are produced if 2.50  102 g of nitroglycerin explodes?
4. Sodium carbonate is very important in the manufacture of glass but it is scarce in Nature.
Most of it is made from two more abundant compounds: calcium carbonate and sodium
chloride (calcium chloride is a by-product). How many grams of sodium carbonate can be
made from 255 g of calcium carbonate?
UNIT 3
4
Limiting Reactant Problems
5. In the zinc iodide lab, assume 0.50 g of zinc was combined with 0.50 g of iodine. Which
substance is the limiting reagent? How many grams of zinc iodide can be made from these
quantities?
6. Calcium metal will displace neodymium (Nd) from NdF3. If 26.5 g Ca and 43.2 g NdF3 are
mixed, how many grams of Nd are formed?
UNIT 3
5
7. Sodium metal reacts violently with water to produce sodium hydroxide and hydrogen gas. If 14.0 g of
sodium is added to 1542 g of water, how many grams of sodium hydroxide can be produced?
UNIT 3
6
Name:___________________________ Per.:____
Date:_______________
The Carbonate Project, Part 1
Purpose: to determine the identity of the unknown alkali metal carbonate
Method: In the first part of the experiment I will mass a sample of the compound and a sample of
6 M HNO3. I will slowly add the acid to the compound until there is no further evidence of
reaction or remaining solid. At this point I will re-mass the acid container (now “empty”) and
also the reaction container.
From the data I can determine the mass loss of the reaction mixture which is CO2 that escaped
during the reaction. The moles of CO2 that escaped are equal to the moles of the carbonate since
the formula is M2CO3 (there is one C in each formula unit).
I can find the molar mass of the carbonate by dividing the mass of the original compound used
by the moles from above. Once I know the molar mass, I can find the atomic mass of “M” by
subtracting away the molar mass of CO3 and dividing by 2. This will allow me to identify the
alkali metal.
unknown # _________
Data:
Trial 1
Item
Trial 2
Mass, g
Item
empty flask
empty flask
carbonate
carbonate
25 mL cylinder + acid
25 mL cylinder + acid
final reaction mix
final reaction mix
emptied 25 mL
cylinder
emptied 25 mL
cylinder
UNIT 3
Mass, g
7
Name:___________________________ Per.:____
Date:_______________
The Carbonate Project, Part 2
Purpose: to determine the identity of the unknown alkali metal carbonate
Method: In the second part of the experiment I will mass a sample of the compound, dissolve it
in minimal water, and add some 0.20 M CaCl2. A precipitate of CaCO3 will form. After warming
to digest the precipitate, I will filter the mixture collecting the precipitate on a pre-massed filter
paper [which has my locker number written in PENCIL on it]. The paper and precipitate will be
dried overnight in the oven before a final massing.
From the data I can determine the mass of CaCO3 produced. Converted to moles, this is the same
number of moles as the original carbonate (both contain one CO3 per unit).
I can find the molar mass of the unknown carbonate by dividing the mass of the original
compound used by the moles from above. Once I know the molar mass, I can find the atomic
mass of “M” by subtracting away the molar mass of CO3 and dividing by 2. This will allow me
to identify the alkali metal.
Data:
unknown # _________
Item
Mass, g
carbonate
prepared filter paper
paper + precipitate after drying
UNIT 3
8
LAB: The Carbonate Project
By now you have a good feeling for what a chemical reaction looks like. But in addition to the information
that is contained in the appearances of a chemical process, there is information in the quantities of a chemical
process. In the determination of empirical formulae you have seen that reacting masses are important (and fixed).
The ratios in which substances combine in a chemical reaction may be measured by mass, but it is most convenient
to express them in terms of moles since in that we are "counting" the atoms as they rearrange. Thus in your
determination of the empirical formula for zinc iodide you eventually reduced all of your mass data to a ratio of
moles zinc/moles iodine.
The reaction of an alkali metal carbonate (Na2CO3 for example) with nitric acid is an interesting system to
look at because it is simple and yet versatile. The quantitative relationships (or the stoichiometry) involved in the
reaction can be examined from several different viewpoints. For example, when the acid reacts with the carbonate,
all of the CO32- is converted to CO2. If the gas is allowed to escape, the mass of CO2 lost contains valuable
information about the ratio of HNO3 added to the amount of metal carbonate originally present.
The soluble alkali metal carbonate can also be converted to an insoluble carbonate and the moles of that
solid will be related to the moles of the starting material.
In this experiment we will also introduce the concept of solution stoichiometry, that is, the idea that a
volume of solution can be used to measure the amount of dissolved material added just as a mass measurement can.
This is true, of course, only if the concentration of the solution is known. We conveniently express concentration in
moles/Litre or Molarity (M) because moles are the fundamental measuring units in chemical processes.
Finally, the reaction of HNO3 with the original compound can be followed more closely as an acid/base
titration (carbonate solutions are slightly basic) using an indicator. This final method brings all techniques together:
masses and volumes are used to determine the amount of the metal carbonate.
As described, the planned "experiment" is pretty much another exercise. To make this exercise a little more
sporting, we will add a slight twist. There are five alkali metals, all of which form stable carbonates with the general
formula X2CO3. The latter two (Rb2CO3 and Cs2CO3) are very expensive, but the first three are quite common in
the lab. Your goal in this experiment is to use the data you collect from the three procedures to determine which of
the first three possible carbonates you have been given!
Note that in each procedure as generally described you can determine the moles of CO32- in your sample.
According to the formula, the moles of X are in a 2:1 ratio with this value. The mass of that number of moles of X
can be determined by difference if you know the mass of CO32- lost or reacted. And the mass/mole of X should agree
with the atomic mass of one of the first three alkali metals.
Preparing to experiment
You will be provided with the following materials:
1.
2.
3.
4.
5.
6.
7.
8.
9.
UNIT 3
an unknown alkali metal carbonate (use about 2 g in the first procedure, about 1 g in the second and
about 0.10 g in the third)
6 M HNO3 (that's 6 moles HNO3/Litre) (use about 25 mL in the first procedure)
0.20 M CaCl2 solution (use about 80 mL in the second procedure)
0.10 M HNO3 (use for the titration)
methyl orange/indigo carmine mixed indicator (use about 5 drops)
filter paper circles
a 600 mL beaker
a 50 mL buret
a magnetic stirrer
9
Design an experiment in which the mass loss of a sample of alkali metal carbonate is determined after reaction with
acid in a 250 mL Erlenmeyer flask. Remember that an immediate goal is to limit the loss of material, and spray will
be generated when the acid comes in contact with the solid.
Design an experiment by which you can determine the mass of calcium carbonate formed when a solution of alkali
metal carbonate (use up to 300 mL of distilled water to get your sample to dissolve) is mixed with a solution of
calcium chloride.
Design an experiment by which you can accurately measure the amount of HNO3 needed to "neutralize" a sample of
an alkali metal carbonate using an indicator.
Adapted from: A Carbonate Project Introducing Students to the Chemistry Lab, Emily Dudek, J. Chem. Educ., 1991,
v. 68, p. 948
Pre-lab take-home quiz
Answer these questions on a separate sheet of paper to be turned in on the day you do this experiment. Note that
these three questions are based on the three different experimental techniques. Because of scheduling, you
may not do the experiment in the same order in which it is written. Each night you should only do the
question below that is pertinent to the work you will do in the lab on the following day.
1.
A sample of 1.60 g of X2CO3 is dissolved in excess nitric acid and all of the resulting CO2 expelled. The added
mass of acid is 26.10 g. The final mass of the solution mixture is 26.75 g.
a. How many grams of CO2 were lost?
b. How many moles of CO2 is this?
c. How many moles of X2CO3 must have been present (each X2CO3 makes one CO2)
d. Based on your answer to (c) and the starting mass of the sample, what is the approximate molar mass of
unknown carbonate?
e. Subtracting the molar mass of CO32- from the answer to (d) and dividing by 2 should give an approximate
atomic mass for X. Do this and SHOW WHY THIS IS TRUE.
2.
A 0.75 g sample of X2CO3 is dissolved in water and 80 mL of 0.20 M CaCl2 solution is added.
a. Write a balanced molecular equation for the reaction that takes place.
b. If the mass of dried, washed precipitate is 0.71 g, calculate the moles of CaCO3 formed.
c. Based on the ratio in your balanced equation, how many moles of X2CO3 must have reacted? How many
moles of CO32- is this?
d. Find the approximate molar mass of X2CO3 from the initial sample mass and the moles in (c).
e. Perform an operation similar to that described in (1.e) to get the approximate atomic mass of X.
f. As an afterthought, calculate the moles of CaCl2 in the 80.0 mL of 0.20 MCaCl2. How many moles of Ca2+
does this provide?
g. Show that the amount of CaCl2 in (f) provides more Ca2+ than is needed to react with all the CO32- present
and form a precipitate
3.
The reaction of HNO3 with X2CO3 produces CO2, water, and XNO3.
a. Write a balanced molecular reaction for this process.
b. If 17.4 mL of 0.10 M HNO3 was required to "neutralize" a 0.12 g sample of X2CO3, how many moles of
HNO3 is this?
c. Based on your balanced equation, how many moles of X2CO3 must have reacted with the moles of HNO3 in
(b)?
d. Use the sample mass and moles in (c) to calculate an approximate molar mass for X2CO3.
e. Determine an approximate atomic mass for X using this information (see 1.e)
UNIT 3
10
Technique
1.
Digesting a precipitate
Frequently precipitates form in solution as very fine crystals that will pass through filter paper, yielding poor results.
Such is the case with calcium carbonate. This effect can be minimized by "digesting" the precipitate, helping larger
crystals to form. Usually this is done by heating the mixture (in this case, just to boiling) and then allowing it to cool
somewhat before filtering. If the digestion is complete and the filter paper is prepared properly, the filtrate should be
clear and free of solid.
2.
Filtering
Gravity filtration is a somewhat slow process for separating precipitates from mixtures, particularly those which
tend to remain suspended. A circular piece of paper is folded as shown in the diagrams below and then opened into a
cone. This cone is fitted into a funnel. A container under the funnel catches the liquid (called the filtrate) and the
solid (the residue) is caught in the paper.
For quantitative work, the paper cone should be prepared and massed before fitting it into the funnel. Then the cone
is held in place in the funnel by adding a small amount of solvent (generally water) with an eyedropper or wash
bottle. It should fit smoothly on all sides.
3.
Titration
You already know the essentials of titration. But the buret is the commonly used volumetric glassware for this
procedure. It is graduated in 0.1 mL intervals, like a long graduated cylinder, but more accurate. It has a valve at the
bottom for letting the solution out in measured amounts, from a slow drip to a steady stream.
The solution to be titrated is placed in a beaker or flask under the filled buret. Indicator is generally added. The
mixture should be stirred as the titrant is added. This can be done by hand or a device called a magnetic stirrer can
be used. This is a motor with a magnet attached to its shaft mounted under a platform . A small teflon-coated magnet
is placed inside the beaker or flask and when the motor turns, the magnet follows it. Mixing insures that the reaction
will be complete as the titrant is added.
At the beginning of a titration you can generally go pretty fast, but you should slow down when you notice a color
change occurring around where the titrant is dripping in, finally going drop-by-drop to get the exact point at which
the color changes. In this experiment, the indicator changes from green to purple.
A buret should always be rinsed thoroughly after use and clamped upside down with the valve open so that the
entire body will drain.
UNIT 3
11
4.
Massing for transfer
Whenever substances must be quantitatively transferred (i.e., transferred without loss) you need to think carefully
about how you measure. The general rule is to choose a procedure that minimizes transfer. For example, in the first
method when you need to know the total mass of the mixture before reaction and the total mass after reaction, you
should be careful how you measure the added HNO3, remembering that not all of it will pour out of the graduated
cylinder and into the flask (some will cling to the sides of the cylinder). Thinking about these kinds of operations
before you do them will ensure that you record sufficient data so that you will be able to interpret your results later.
The chemicals
Lithium carbonate, sodium carbonate and potassium carbonate are the possible compounds you might have in
this experiment. All are relatively hygroscopic, but K2CO3 is the worst. So they are stored in the oven or a desiccator
when not in use. Lithium carbonate is used in pottery glazes and for treatment of manic psychosis. Dry sodium
carbonate, sometimes called soda ash, is a white powder with a bitter taste. The hydrated form is sometimes called
sal soda or "washing soda" and is still used in some laundry detergents and as a general cleanser. Potassium
carbonate is used in the manufacture of some soaps and glasses as well as for tanning leather.
Calcium chloride is obtained as a by-product in the manufacture of sodium carbonate (The Solvay process). It is
very hygroscopic and the anhydrous form is used as a drying agent. It is also useful for fireproofing fabrics, for
melting ice and snow on the ground and roads, in concrete mixtures for greater strength, and as a brine for filling
inflatable tires on tractors to provide better traction.
Methyl orange is a complex hydrocarbon which is used as an indicator for combinations of weak bases and strong
acids. It changes from red to yellow as the acidity of a solution decreases.
Indigo carmine is another complex hydrocarbon which is sensitive to the amount of acid or base in a mixture. It has
been used as a food dye (FD & C Blue #2) as well as a general purpose dye. It is yellow in strongly basic solutions
and blue in less basic solutions.
Analysis
1. Use your data from the mass loss of CO2 to determine a value for the atomic mass of the metal X in your
carbonate unknown and make a tentative identification of the metal. Discuss any variation in the mass you calculate
and the known atomic mass of the metal you choose (e.g., is your calculated molar mass higher or lower than the
actual molar mass and why). You must calculate the atomic mass for both trials.
2. Use your data from the precipitation of CaCO3 to determine a second value for the atomic mass of X. Discuss any
variation in the mass you calculate and the known atomic mass of the metal you choose (e.g., is your calculated
molar mass higher or lower than the actual molar mass and why).
3. Use your titration data to determine a third value for the atomic mass of X. Discuss any variation in the mass you
calculate and the known atomic mass of the metal you choose (e.g., is your calculated molar mass higher or lower
than the actual molar mass and why).
4. When you have completed all three parts and the flame test (next page) decide which unknown carbonate you had
and explain your choice briefly.
UNIT 3
12
An afterthought:
The alkali metals are known for the brilliant and distinctive colors they give to a burner flame when their
compounds are injected. According to the Handbook of Chemistry and Physics, these are the colors:
Lithium: deep red or reddish-purple
Sodium: bright yellow or yellow-orange
Potassium: pale lavender
Go back to the lab to check your result!
To do a flame test, dissolve a speck of the unknown in 1 mL or less of 6 M HCl. Dip a wire loop into a small amount
of 6 M HCl and hold it in the flame. Repeat this until there is no color added to the flame before the wire itself
becomes red-hot. Now dip the wire into the solution you made and place it in the flame again. Repeat this several
times until you can identify the color.
Congratulations! You have just completed a very challenging set of experiments!
UNIT 3
13
UNIT 3
14
Unit 3 Sample Test
You will have 5 multiple choice questions, 3 problems, 2 net-ionic reactions, and one essay
question. From this test until the end of the year you will have a choice of two out of four
reactions to complete and write as balanced net-ionic equations. You will also need to identify
the type of reaction. As with the Unit 2 Test, the Activity series will be provided along with the
remaining solubility rules (you need to know #1 and #2)
The following are representative of typical multiple choice questions but do not necessarily
indicate topics to be addressed on your actual test.
_____1.
What is the concentration of nitrate ions in a 0.50 M Ca(NO3)2 solution?
a. 0.50 M
b. 1.0 M
c. 0.25 M
d. 0.125 M
_____2. Which one of the following processes would be used to obtain the elements A and B
from the compound AB?
a. combustion
_____3.
b. distillation
c. precipitation
d. decomposition
c.
d.
A redox reaction takes place when
a. copper is added to FeSO4 solution
b. silver is added to Zn(NO3)2 solution
c. iron is added to AgNO3 solution
d. zinc is added to Mg(NO3)2 solution
e. lead is added to Al2(SO4)3 solution
_____4.
Which object below is an erlenmeyer flask?
a.
_____5.
b.
Which reaction below could be followed by using an acid/base indicator?
a. N2 + 3 H2  2 NH3
b. 2 KOH + H2SO4  K2SO4 + 2 H2O
c. S + O2  SO2
d. Ba2+ + SO42-  BaSO4
e. Zn + Cu2+  Zn2+ + Cu
UNIT 3
15
The next section consists of representative problems which might be found in the problem
section. All students are expected to work on all 3 of the required problems.
6. Magnesium metal, like most metals, reacts vigorously with mineral acids such as
hydrochloric acid in a redox displacement reaction.
a. Write a balanced molecular equation for this reaction:
______________________________________________________________________________
b. How many grams of magnesium metal will be consumed by 48 g of hydrochloric acid,
assuming an excess of magnesium?
c. How many grams of hydrogen would be expected to form in the reaction if 48 g of
hydrochloric acid is combined with 8.0 g of magnesium? What is the limiting reagent?
7. Some of the substances commonly used in over-the-counter preparations for neutralizing
excess stomach acid include Mg(OH)2 and Al(OH)3.
a. Write a balanced molecular reaction for the neutralization of each by hydrochloric acid.
______________________________________________________________________________
______________________________________________________________________________
b. When 0.25 g of ONE of the compounds above is titrated with 0.10 M HCl, 96.2 mL are
required to reach the endpoint. Which substance is it?
UNIT 3
16
8. A 100.0 mL sample of 0.200 M potassium hydroxide solution is mixed with 100.0 mL of
0.200 M magnesium nitrate solution.
a. Write a balanced molecular equation for the reaction which occurs:
______________________________________________________________________________
b. What is the precipitate that forms?____________________
c. How many grams of the precipitate would be formed?
9. A 10.00 mL sample of sulfuric acid from an automobile battery requires 32.47 mL of 2.15 M
sodium hydroxide solution for complete neutralization. What is the Molarity of the sulfuric acid?
(hint: it might be helpful to write out the reaction first)
UNIT 3
17
The next section consists of representative reactions to complete and write balanced net-ionic
equations for. Note that some reactions do not occur in aqueous solution and thus molecular
equations are all that would be needed. Each student is expected to choose two from this section.
In addition to a periodic table, solubility rules 3-6 and a copy of the activity series will be
included with the test. Mixtures which result in no net change need not be completed. Simply
write "NO REACTION". Non-trivial redox reactions are indicated with * and only the balanced
net-ionic equation is required.
10. For each of the following, complete the word equation, write a balanced net-ionic reaction
and tell what type of reaction it is (precipitation, acid/base, acid/metal oxide, acid/carbonate,
redox, etc.). Unless otherwise noted, all reactions occur in aqueous solution.
a. lead(II) nitrate + sodium iodide  ______________________________________________
______________________________________________________________________________
______________________________________________________________________________
type:__________________
b. copper(II) carbonate + hydrochloric acid  _______________________________________
______________________________________________________________________________
______________________________________________________________________________
type:__________________
c. silver metal + copper(II)  ____________________________________________________
______________________________________________________________________________
______________________________________________________________________________
type:__________________
*d. an acidic solution is prepared containing chromate ions and iron(II) ions; among the products
after reaction are iron(III) ions and chromium(III) ions
______________________________________________________________________________
______________________________________________________________________________
The final section of the test will consist of one essay question selected from the following topics:
--technique involving balances
--technique involving filtering
--lab equipment (glassware, etc.)
--use of the activity series
UNIT 3
18