# Stoichiometry Reactions: Ammonia

## What is stoichiometry?

Stoichiometry is a method to figure out how much/many of the things are possible based on the balanced equation when given certain amount of something in the equation. For example, 2 graham crackers, a marshmallow, and a chocolate square makes a s'more. How many s'mores can you make with 6 graham crakers, 5 marshmallows, and 4 chocolate squares? The answer is 3 because after 6 graham crackers, there won't be any left. Stoichiometry is actually really easy and it is more easier than it seems.

## Terms

Excess Reactant: the reactant that has higher yield than the limiting reactant

Limitng Reactant: the reactant that has the lowest theoretical yield due to its lack of quantity. This is the reactant that would be exhausted the quickest

Molar mass: the amount of grams it would weigh for the one mole of an object

Mole: a standard scientific unit for measuring large quantities of very small entities such as atoms, molecules, or other specified particles

Product: the substances that result form the recombination of atoms

Reactant: the substance that takes part in and undergoes change during a reaction

Stoichiometry: the measurement and calculation of the amounts of reactants and products in chemical equations

Theoretical yield: how much something can be produced when given value of something else; equivalent to the limiting reactant's yield

Percent yield: displays how much of a reaction was produced versus the predicted, or theorized amount. It is displayed in percentages and it is actual yield/ theoretical yield

## How to Solve Stoichiometry Problems

For our purposes, we are gonna use nitrogen and hydrogen as an example.

• N2 + 3H2 -> 2NH3 H A
This reaction is synthesis because two substances are forming to create one substance.

After balancing the equation out, we can now start looking at the stoichiometry.

1. Balancing equation: you have to add coefficients in the equation to make sure that there remains the same number of substances remain. This is to fulfill the law of conservation of mass. The equation above is already balanced.
2. Molar mass: molar mass is obtained from finding its each respective molar masses of every element listed in the substance and multiplying by its subscripts and adding the product of all the elements in the substance. You do not multiply by its coefficient. In this case, 2*(14.007)= 28.014 g for the nitrogen, 2*(1.008)= 2.016 g for hydrogen, and 14.007+3*(1.008)= 17.031 g for ammonia.
3. Mole to mole conversions: in order for mole to mole conversion, you must understand that the coefficients on the equation actually tells you the ratio of the moles between substances in the equation. Because of this, the product of the reactant A's given mole and reactant B's coefficient should equal to the product of the reactant B's given mole and reactant B's coefficient. Much more simply, multiply the given mole of reactant A by the ratio of the two substances, where reactant B's coefficient on the numerator and the reactant A coefficient on the denominator. For this example, I will use 10.04 moles of nitrogen to find the number of moles of ammonia. I will multiply 10.04 by 2 as the ratio between the ammonia and the nitrogen is 2:1. As a result, I get 20.08 moles of ammonia.
4. Mass to mass conversions: mass to mass is one more step than mole to mole conversions. First you have to find how many moles of the given substance you have when given its mass. This is achieved by diving the given mass by its molar mass. Then, you must multiply by its ratio in order to find out how many moles of substance B it can produce (this is mole to mole conversion). Then, additionally, you have to multiply the amount of mole of reactant B by its molar mass to get the final mass. This is to actually calculate how much mass it is when its not 1 mole of that final product. For this example, I will use 12.1 g of nitrogen and convert it to mass of ammonia. I will divide 12.1 by 28.014, multiply the quotient by 2, and multiply the product by 17.031. The solution comes out to 14.71229385 and this is reduced to 14.7 g as it was rounded to least significant figure in this arithmetic.
5. Limiting and excess reactant: because limiting reactant limits the amount of product it can be produced due to its lack of quantity, the lowest yield out of the two reactants is the limiting reactant. Conversely, the other reactant that does not have the lowest yield is the excess reactant. For this, you must solve mass to mass conversions for every single reactant given and compare the results. In this example, I will be using 12.1 g for nitrogen and 20.15 g of hydrogen to see how much mass they can produce ammonia based on their given mass. For nitrogen, we already have calculated the mass of ammonia, which was 14.7 g. For hydrogen, you divide 20.15 by 2.015, multiply the quotient by 2/3 (ratio), and multiply the product by 17.031. This produces 113.4836806, or 113.5 g of ammonia when given 20.15 g hydrogen. The limiting reactant in this case is nitrogen as it only allows to create 14.7 g than 113.5 g of hydrogen's; the excess reactant is the hydrogen.
6. Theoretical yield: the theoretical yield is the yield of the limiting reactant, which is 14.7 g.
7. Percent yield: its the quotient of the actual amount and the theoretical amount. In this case, assume that 6.5 g of ammonia has been produced with the same conditions as before when calculating the theoretical yield. 6.5 divided by 14.7 is approximately .44 and you have to multiply by 100 to get the percentage. Thus, the percent yield is 44%.

## Research About the Reaction: Ammonia

Ammonia is the product of two nonmetals called nitrogen and hydrogen. It is colorless, has a highly irritating odor, and it is a gas at a room temperature. It dissolves easily in water to form ammonium hydroxide; it is also easily compressed to form a clear liquid under pressure. Ammonia is not highly flammable although containers of ammonia may explode when exposed to high heat. Ammonia is also one of the most produced industrial chemicals in United States. It has both industry and commerce usage, and it is also found in natures, in humans and the environment. About 80% of its production is utilized as a fertilizer in agriculture. It is also used as a refrigerant gas and the manufacture of plastics, explosives, textiles, pesticides, dyes, and other chemicals. For humans, this product is harmful when exposed to large amounts. Ammonia will create ammonium hydroxide at the contact of the moisture from skin, eyes, oral cavity, respiratory tract; it causes cellular destruction and inflammation. Ingestion of ammonia will create corrosive damage to the mouth, throat and the stomach. In the nature, ammonia takes part in the nitrogen cycle and is produced in soil from the bacterial processes.

## Images and Sources for Additional Help

Balancing Chemical Equations - Chemistry Tutorial

## Sources

The Facts About Ammonia. (n.d.). Retrieved December 10, 2015, from https://www.health.ny.gov/environmental/emergency/chemical_terrorism/ammonia_tech.htm