Section 1: Understanding Chemical Reactions
Consider the burning of propane in a gas grill.
C3H8 (g) + 5 O2 (g)  3 CO2 (g) + 4 H2O (g)
1. Set up the atom cards for the reactants. Use the cards to show how this
reaction progresses.
2. Compare the total mass of the products to the total mass of the reactants.
Explain your reasoning in terms of the atom cards.
3. Compare the number of atoms of each type of element before and after the
reaction. Explain why this must be true in terms of the atom cards.
4. Looking at the carbon dioxide in the balanced equation, explain what the 3
in front of the CO2 means as well as what the subscript 2 tell you. How are
these numbers different?
5. For each statement below, indicate whether it is possible or impossible
during a chemical reaction.
a. Atoms can be rearranged to form new products.
b. Bonds in the reactant molecules can be broken.
c. An atom can change the type of element it is during the reaction. For
example, under the right reaction conditions a hydrogen atom can turn
into a carbon atom.
Section 2: Balancing Chemical Equations
1. Use the atom cards to make the chromium(III) carbonate. Get your
compound checked off before continuing.
Cr2(CO3)3
a. Draw Cr2(CO3)3 in the space to the right.
2. What does the subscript 2 after the Fe mean?
3. What does the subscript 3 after the O mean?
4. What does the subscript 3 after the parenthesis mean?
5. Use the atom cards to balance the following reaction. Show your teacher
how you did this process.
___ Cr2(CO3)3 (s) ___ Cr (s) +___ CO2 (g) +___ O2 (g)
6. Balance the following reactions without atom cards:
____ AlPO4 (aq) + ____ MgCO3 (s)  ____ Mg3(PO4)2 (aq) + ____ Al2(CO3)3 (s)
____ HNO3 (aq)+ ____ Cu (aq)  ____ Cu(NO3)2 (s)+ ____ NO2 (g)+ ____ H2O (l)
Section 3: Using a Balanced Equation to Make Predictions
Consider the reaction between carbon and hydrogen to produce ethane gas.
2 C (s) + 4 H2 (g)  C2H8 (g)
1. Use words to describe what is happening in this reactions.
2. How many hydrogen molecules would you need to have to use up 8 carbon
atoms? Show your work using the fraction method.
3. How many ethane molecules could you produce if you used up 16 hydrogen
molecules? Show your work using the fraction method.
4. How many moles of hydrogen would need in order to produce 0.375 moles
of ethane? Show your work using the fraction method.
5. How many moles of carbon would you need to use 13.68 moles of
hydrogen? Show your work using the fraction method.
6. David completes the following two calculations. What are the units and
physical meanings of each of David’s answer?
12.8𝑚𝑜𝑙𝑒𝐻2 1𝑚𝑜𝑙𝑒𝐶2𝐻8
(
) = 3.2
1
4𝑚𝑜𝑙𝑒𝑠𝐻2
7.15𝑚𝑜𝑙𝑒𝐶 4𝑚𝑜𝑙𝑒𝐻2
(
) = 14.3
1
2𝑚𝑜𝑙𝑒𝑠𝐶
Section 4: Understanding and Calculating Molar Mass
1. If you look on the periodic table the molar mass (atomic mass) of carbon is
12.011 g/mol. Explain what this number tells you about carbon.
2. Use the atom cards to build the following compound. Get your compound
checked by your teacher and then draw it in the space below.
Cr2(SO4)3
3. Describe what the subscript
2 means in the formula.
4. Describe what the subscript
4 means in the formula.
5. Describe what the subscript 3 means in the formula.
6. Consider the compound aluminum ferrocyanide, Al4(Fe(CN)6)3. How many
atoms of each present are in one aluminum ferrocyanide?
a. Al _____
b. Fe _____
c. C _____
d. N _____
7. Calculate the molar mass of aluminum ferrocyanide. Show your work.
8. Calculate the molar mass of the flowing compounds. Show your work.
a. H2SO4
c. Pb(IO3)4
b. C5H9(NH3)2COOH
Section 5: Converting Between Moles and Grams
1. The molar mass of a methane, CH4, is about 16 g/mol. What does this
number tell you about methane?
a. If you have 16 g of methane you have ______ moles of methane, which
is the same as having ________________ molecules of methane.
b. If you have 32 g of methane you have _______ moles.
c. If you have sample that contains 22.8g of methane you have
( less / more ) than a mole of methane. Explain your reasoning below.
2. The molar mass of calcium phosphate, Ca3(PO4)2, is about 310 g/mol. If you
have 0.310 g of calcium phosphate, you have ( less / more ) than a mole of
methane. Explain your reasoning below.
3. The molar mass of a water is about 18 g/mol. If you 1.5 moles of a
substance, you will have to measure ( less / more ) than 18g of water.
Explain your reasoning below.
4. The molar mass of carbon dioxide is about 44 g/mol. If you need 0.15moles
of carbon dioxide, you will need ( less / more ) than 44g of carbon dioxide.
Explain your reasoning below.
5. Use the fraction method to calculate the number of moles of hydrogen
peroxide, H2O2 in 16.57g. Show your work.
6. Use the fraction method to calculate the number of grams of dinitrogen
pentaoxide, N2O5 in 0.156 mol. Show your work.