CHEMISTRY 161 - EXPERIMENT 6
REACTIONS & STOICHIOMETRY: A Study in How Chemicals Interact
INTRODUCTION:
PART I:
If you examine a car that has been in a junkyard for a while, you will notice that it has rusted. Rusting is a
slow chemical reaction of the iron in a car’s body with oxygen gas. If a piece of sodium is put into water, a much
more rapid chemical reaction occurs. Sodium reacts with water to produce sodium hydroxide and hydrogen gas.
During this reaction, enough heat is liberated to ignite the hydrogen gas causing it to explode.
Chemists observe what is happening in a chemical reaction and try to describe it in language that is simple
and clear. A chemical equation uses formulas and symbols to describe the substances involved in a reaction, the
physical state of the substance, the use of a catalyst, and relative proportions. The general form of n equation is:
Reactants  Products
In this investigation you will perform a series of reactions and make careful observations of the changes that
occur. Using simple tests and your knowledge of chemistry, you will determine the identity of the products. With
this information, you will write chemical equations to describe the reactions. In this lab you will use chemical
reactions to describe what happens in chemical reactions.
PART II:
In chemistry, the limiting reagent, also known as the "limiting reactant", is the chemical that determines how
far the reaction will go before the chemical in question gets used up, causing the reaction to stop. The chemical
of which there are fewer moles than the proportion requires is the limiting reagent.
Example
Consider the combustion of benzene:
1.5 mol C6H6 ×
This means that 11.25 mol O2 is required to react with 1.5 mol C6H6. If only 7 mol O2 is present, the oxygen will be
consumed before benzene. Therefore, O2 must be the limiting reagent.
This conclusion can be verified by comparing the mole ratio of O2 and C6H6 required by the balanced equation with
the mole ratio actually present:
required:
=
actual:
=
Since the actual ratio is too small, O2 is the limiting reagent.
Consider a typical thermite reaction:
If 20.0 g of Fe2O3 are reacted with 8.00 g Al(s) in the thermite reaction, Which reactant is limiting?.
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First, determine how many moles of Fe(l) can be produced from either reactant.
Moles produced of Fe from reactant Fe2O3
Moles produced of Fe from reactant Al
Because the moles Fe produced from Fe2O3(0.254mol) is less than the moles Fe produced from
Al(0.297mol), Fe2O3 is the limiting reagent.
By looking at chemical equation for the thermite reaction, the limiting reagent can be found based on the ratio of
moles of one reactant to another and the total atomic mass of the reactant compounds.
PROCEDURE:
PART I:
1. For each of the reactions, record in the Data Table observations such as the appearance of the reactants;
evidence that a chemical reaction has taken place; the results of tests performed on any gases produced; the
appearance of the products; and any other relevant data.
2. Obtain a piece of magnesium ribbon. Light the lab burner. Holding the magnesium with your tongs,
carefully place it in the flame. Once ignited, hold the burning magnesium over a watch glass to catch any
debris. Place any unburned magnesium on the watch glass.
3. Obtain a piece of copper foil. Light the lab burner. Holding the copper with your tongs, carefully place it in
the flame and heat it until it is red hot. Remove it from the flame and allow it to cool. Scratch the surface
of the metal with a sharp object.
4. Place a test tube in a test tube rack. Have a second test tube ready in a test tube holder. Add 5-10mL of
3.0M Hydrochloric acid to the first test tube. Drop a 2cm piece of magnesium ribbon into the acid and
immediately invert the second test tube over the first to catch any gas released. When the reaction appears
to have ended, ignite a wooden splint and quickly test the collected gas for flammability by holding the
burning wood splint near the mouth of the second test tube.
5. Carefully place about 1 spatula of (NH4)2CO3 into a test tube. Holding the test tube with a test tube holder,
heat the solid by holding the test tube in the flame for a few seconds, then removing it for a few seconds.
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As you heat the solid, carefully waft the air toward your nose to detect any odor. Continuing to heat the
solid, place a burning splint at the mouth of the test tube. Finally, as heating continues, place a piece of blue
cobalt paper just inside the mouth of the test tube. Put the test tube in the rack to cool and turn off your
burner.
6. Place approximately 10-15mL of hydrogen peroxide, H2O2, into a test tube. Have a wooden splint and
matches ready. Add a very small amount of MnO2 to the hydrogen peroxide. As the reaction occurs, light
the splint and allow it to burn freely for about 5 seconds. Blow the flame out and place the glowing splint
halfway into the test tube.
7. Place a drop of potassium iodide solution, KI, on a watch glass. Add a drop of lead (II) nitrate solution,
Pb(NO3)2.
PART II:
1. Mass together a clean, dry evaporating dish and a watch glass. Record the mass to the thousandths place.
2. Place approximately 2 grams (but no more than 3.5 grams) of NaHCO3 into the evaporating dish. Reweigh
the dish, NaHCO3, and the watch glass to the thousandths place. Calculate the exact mass of NaHCO3 in
the dish.
3. SLOWLY add 7.0 mL of 6.0 M HCl to the NaHCO3 by pouring the HCl carefully into the evaporating dish.
To avoid splattering materials out of the dish, leave the watch glass on top of the dish and add the
HCl through the lip of the evaporating dish. You may want to do this in small portions to avoid excess
splattering. If any of the starting materials get splattered out of the dish, you will need to start over.
4. Holding the watch glass on top of the evaporating dish, gently swirl the contents of the dish together so that
all of the NaHCO3 comes into contact with the HCl. Use minimal distilled water from your wash bottle to
rinse the underside of your watch glass into the reaction mixture in the evaporating dish.
5. Set the covered evaporating dish on a hot plate or a tripod set-up with a Bunsen burner. The contents of
the dish need to be at a gentle boil so that the water can evaporate slowly. The NaCl needs to be completely
dry before removing the dish from the heat source. There should be no signs of moisture (no water
droplets or traces of steam!) in the solid or on the underside of the watch glass.
6. Allow the evaporating dish to cool to room temperature (a minimum of 10 minutes - most likely longer).
Then mass the evaporating dish, its contents (NaCl) and the watch glass.
7. Heat the evaporating dish for a few more minutes, allow to it cool and then re-mass. If the same mass is
obtained, then you can assume all of the w ater was driven off. If not, this procedure should be repeated
until a constant mass is obtained.
POST-LAB QUESTIONS:
PART I:
1. Write a balanced chemical reaction for each of the reactions performed. Include the physical state of each
substance.
2. Classify each reaction as synthesis, decomposition, single replacement or double replacement.
3. A positive test for hydrogen was the “pop” test. What chemical reaction was occurring? Write a balanced
equation representing the reaction.
PART II:
1.
2.
3.
4.
Which reactant is the limiting reagent?
How many moles of NaCl do you think you can make based on this knowledge (theoretical yield)?
Convert moles of NaCl to grams of NaCl.
Find the average percent yield of NaCl.
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Name
RESULTS SHEET
REACTIONS & STOICHIOMETRY: A Study in How Chemicals Interact
PART I:
Reaction
Observations
Burning Mg
Mg and HCl
Heating Cu
Heating (NH4)2CO3
H2O2 and MnO2
KI and Pb(NO3)2
PART II:
Trial 1
Mass of dry evaporating dish and watch glass, g
Mass of dry evaporating dish, watch glass and NaHCO3, g
Mass of NaHCO3, g
Moles of NaHCO3, moles
Volume of 6.0M HCl added, mL
Moles of HCl, moles
Mass of evaporating dish, watch glass and products (after cooling) , g
Mass of evaporating dish, watch glass and products (after reheating and cooling) , g
Mass of evaporating dish, watch glass and products, g
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