Fall ‘17
CHE 2A Content Review
Week 2- October 16th & 18th
CHE 2A Content Review #2
Topics:
● Limiting Reagents
Limiting Reagents Pt. I
Limiting Reagent (LR): The reactant that is completely used up in a chemical
reaction
Excess Reagent (ER): The reactant that is in excess amount after a chemical reaction
When you don’t have enough of one reactant, the whole chemical reaction will stop after
using up that limited amount of reactant. In other words, the limiting reagent will limit the
amount of product produced. In the same reaction, the other reactant will be the excess
reagent since there will be more of it to run the reaction.
Example: Building a bike with wheels & handlebars
+ =
“Balanced Chemical Reaction”: 1 Handlebar + 2 Wheels = 1 Bike
In order to produce 1 bike, we need 1 handlebar and 2 wheels. If you had 3
handlebars and 4 wheels, how many bikes can you produce?
Answer using stoichiometry:
1. Find out how many bikes you can produce from each reactant (your wheels and
handebar)
Handlebar: 3 handlebars(1 bike/1 handlebar) = 3 bikes
Wheels: 4 wheels(1 bike/2 wheels) = 2 bikes
2. Compare the amount of product produced from each reactant
a. The reactant the produces the least amount of product will be your limiting
reagent
Because your 4 wheels produce the least amount of bikes, “wheels” will be your limiting
reagents and your “handlebars” will be your excess reagent. This must mean that with 3
handlebars and 4 wheels, you can only produce 2 bikes.
Limiting Reagents Pt. II (Theoretical & Percent Yield)
Theoretical Yield: The expected/calculated amount of product produced from your
limiting reagent
Percent Yield: The amount of product you actually produced in a reaction
- Percent yield % = experimental yield/theoretical yield x 100%
When you calculate the mass or moles of a product from your limiting reagent in some
reaction, you are calculating the amount you could produce. (this is your theoretical
yield) When you run this reaction in a lab, you will likely produce a lesser amount of
product when compared to your theoretical yield. (this is your experimental yield) To
find out how much of your theoretical yield you produced in your experiment, you use
the above formula. (This is your percent yield)
Problem 1: You react 3.25 g of NH3 with 3.5 g of O2. Identify your limiting/excess
reactant and theoretical yield.
Chemical Reaction: NH3 + O2 → NO + H O
2
1. Balance Reaction
4NH3 + 5O2 → 4 NO + 6H O
2
2. Calculate amount (mass) of product from each reactant (if there’s more than one
product, use one throughout your calculations)
NH3 : 3.25g NH3(1mol NH3/17.03g
NH3)(4mol
NO/4 mol NH3)(30.01g NO/1mol NO)= 5 .73g NO (LR)
O2: 3.5g O2(1mol O2/32.00g
O2 )(4mol
NO/5mol O2)(30.01g
NO/1mol NO)= 2 .63g NO (ER)
3. Compare the amounts of product produced from each reactant to identify
limiting/excess reactant and theoretical yield
Because our given amount of O2 produces the least amount of product (NO), O 2 is our limiting
reagent. So, N
H3 must be our excess reagent. Now that you know your limiting reagent, the
amount of product produced from O2 (2.63g NO) is the theoretical yield
4. How much of your excess reagent is left after this reaction?
a. First calculate how much of your excess reagent (NH3) you actually
needed to react with your limiting reagent (start with your theoretical yield)
2.63g NO(1mol NO/30.01g NO)(4mol NH3/4mol NO)(17.03g/1mol NH3 )= 1.49g NH3
b. Subtract your starting amount of the excess reagent by how much you
actually needed to determine how much is left over
3.25-1.49= 1.76 g NH3
5. If you ran this reaction and obtained 2.25g NO, what is your percent yield?
[percent yield % = (experimental yield/theoretical yield) x 100%]
Percent yield % of NO = (2.25g NO/2.63g NO) x 100% = 85.6%
