3. Be able to perform a chemical reaction on any scale using the mole. This is perhaps the most common use of this handy topic.
Ex. To make 2 moles of water from the elements, the balanced chemical equation is: _________ + __________ --> _______ so I should mix ___g of hydrogen and ___g of oxygen
4. For any chemical reaction, be able to determine the number of molecules that are involved.
Ex. The above reaction uses ____molecules of hydrogen and ____molecules of oxygen to make ____ molecules of water.
5. Be able to predict the percent composition of any substance.
Ex. Ammonia (NH3): ___%N, ___%H
6. We should be familiar with all types of mole conversions: mole-mole (1 step), mole-gram or gram-mole (2 steps), and mole-mole (3 steps).
Examples: The combustion of one mole of hydrogen with excess oxygen will produce ____ moles of water.
The combustion of 4 moles of hydrogen with excess oxygen will produce _____ grams of water.
The combustion of 4 grams of hydrogen with excess oxygen will produce ___ grams of water. 7. We should be familiar with the concept of a limiting reactant. This is indispensable whenever any amounts of reactants are combined. Unless the stoichiometric amounts are used, this will create a situation where one reactant is in excess, and the other is limiting. We used a non-intuitive but useful method for solving these problems: we find out how much product each reactant will produce, and the lower amount “wins” – it defines the limiting reactant and therefore indicates how much product will form.
For example, what will happen when one gram of hydrogen is combined with one gram of oxygen? 1 g H2 x ____________ x___________ x _________ = g H 2O 1 g O2 x ____________ x ___________ x ___________ = g H2O The limiting reactant is _____, and this reaction will produce ______ g H2O. We learned a quick method for determining how much excess reactant there is- the ratio for the grams of product formed shows how much reactant was consumed. For the reaction above, since oxygen is the limiting reactant, the amount of hydrogen left is:
1 g H2 – 1g H2 (__/___) = _____ g excess hydrogen.
8. Finally, we learned to calculate the yield of a reaction; this is the actual yield/the theoretical yield x 100.
For example, if this reaction above produced 0.1 grams of water, the yield would be ( ___g/____g) x 100 = ____% yield.
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