How To Find Theoretical Yield: A Simple Guide

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Table of contents

  1. Empty
  2. 1. Balance Your Chemical Equation
  3. 2. For Each Reactant, Determine the Molar Mass
  4. 3. Use Molar Mass to Convert the Given Amount of the Reactants to Moles
  5. 4. Calculate the Molar Ratio between the Reactants
  6. 5. Find the Reaction's Ideal Ratio
  7. 6. Pinpoint the Limiting Reactant
  8. 7. Choose the Desired Product and Determine its Ratio to the Limiting Reactant
  9. 8. Multiply the Ratio by the number of Moles of the Limiting Reactant
  10. 9. Use the Molar Mass of the Product to Convert the Result to Grams
  11. Conclusion

Empty

{"blocks":[{"key":"frneg","text":"In a chemical reaction, different agents interact with each other to form a new compound(s). The technical term used to describe the reacting agents is “reactants”, while the resulting compounds are called “products”.","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":217,"style":"color-rgb(0,0,0)"},{"offset":0,"length":217,"style":"bgcolor-transparent"},{"offset":0,"length":217,"style":"fontsize-11pt"},{"offset":0,"length":217,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"dk0lg","text":"The term “theoretical yield” in chemistry refers to the maximum quantity of products a chemical reaction can be expected to form. Determining the theoretical yield of a chemical reaction requires measuring the quantity of the reactants.","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":236,"style":"color-rgb(0,0,0)"},{"offset":0,"length":236,"style":"bgcolor-transparent"},{"offset":0,"length":236,"style":"fontsize-11pt"},{"offset":0,"length":236,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"5m24","text":"This is a simple step-by-step guide on how to find theoretical yield. So, keep reading if you’re interested in understanding these calculations.\n ","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":144,"style":"color-rgb(0,0,0)"},{"offset":0,"length":144,"style":"bgcolor-transparent"},{"offset":0,"length":144,"style":"fontsize-11pt"},{"offset":0,"length":144,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}}],"entityMap":{}}

1. Balance Your Chemical Equation

{"blocks":[{"key":"879ak","text":"Before you start any calculation, you need to make sure that your work foundation is proper. In the case of a chemical reaction, you should balance it before all else.","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":167,"style":"color-rgb(0,0,0)"},{"offset":0,"length":167,"style":"bgcolor-transparent"},{"offset":0,"length":167,"style":"fontsize-11pt"},{"offset":0,"length":167,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"f97kv","text":"This step can be easier if you think of the reaction as a recipe where the number of atoms going in to react with one another must be equal to the number of atoms resulting in the product.","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":188,"style":"color-rgb(0,0,0)"},{"offset":0,"length":188,"style":"bgcolor-transparent"},{"offset":0,"length":188,"style":"fontsize-11pt"},{"offset":0,"length":188,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"dtrum","text":"For example, observe the following chemical reaction between hydrogen gas and oxygen gas to produce water:","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":106,"style":"color-rgb(0,0,0)"},{"offset":0,"length":106,"style":"bgcolor-transparent"},{"offset":0,"length":106,"style":"fontsize-11pt"},{"offset":0,"length":106,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"9do77","text":"H2(g) + O2(g) → H2O(l)","type":"unordered-list-item","depth":0,"inlineStyleRanges":[{"offset":0,"length":22,"style":"color-rgb(0,0,0)"},{"offset":0,"length":22,"style":"bgcolor-transparent"},{"offset":0,"length":22,"style":"fontsize-11pt"},{"offset":0,"length":22,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"4m33o","text":"As you can tell, 2 atoms of hydrogen combine with 2 atoms of oxygen to yield one molecule of water. But, a molecule of water only contains one atom of oxygen, so you need to balance the equation to make sure that the total number of atoms of the reactants equals the total number of atoms of the products.","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":305,"style":"color-rgb(0,0,0)"},{"offset":0,"length":305,"style":"bgcolor-transparent"},{"offset":0,"length":305,"style":"fontsize-11pt"},{"offset":0,"length":305,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"5kqos","text":"For this chemical reaction, you should double the number of water molecules as well as the hydrogen atoms. So, the balanced equation will look like this:","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":153,"style":"color-rgb(0,0,0)"},{"offset":0,"length":153,"style":"bgcolor-transparent"},{"offset":0,"length":153,"style":"fontsize-11pt"},{"offset":0,"length":153,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"1jg5o","text":"2 H2(g) + O2(g) → 2 H2O(l)","type":"unordered-list-item","depth":0,"inlineStyleRanges":[{"offset":0,"length":26,"style":"color-rgb(0,0,0)"},{"offset":0,"length":26,"style":"bgcolor-transparent"},{"offset":0,"length":26,"style":"fontsize-11pt"},{"offset":0,"length":26,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"16pvs","text":"Now, 4 atoms of hydrogen combine with 2 atoms of oxygen to yield 2 molecules of water containing the same number of atoms.","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":122,"style":"color-rgb(0,0,0)"},{"offset":0,"length":122,"style":"bgcolor-transparent"},{"offset":0,"length":122,"style":"fontsize-11pt"},{"offset":0,"length":122,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}}],"entityMap":{}}

2. For Each Reactant, Determine the Molar Mass

{"blocks":[{"key":"9niev","text":"Continuing with the same example, find the molar mass of one molecule of each reactant in the balanced equation. You can refer to the periodic table or any other source.","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":169,"style":"color-rgb(0,0,0)"},{"offset":0,"length":169,"style":"bgcolor-transparent"},{"offset":0,"length":169,"style":"fontsize-11pt"},{"offset":0,"length":169,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"5ed9v","text":"So for the equation 2 H2(g) + O2(g) → 2 H2O(l):","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":47,"style":"color-rgb(0,0,0)"},{"offset":0,"length":47,"style":"bgcolor-transparent"},{"offset":0,"length":47,"style":"fontsize-11pt"},{"offset":0,"length":47,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"dr7v3","text":"The molar mass of a hydrogen molecule is the sum of the molar masses of its 2 atoms. So, 2 × 1 g/mol = 2 g/mol.","type":"unordered-list-item","depth":0,"inlineStyleRanges":[{"offset":0,"length":111,"style":"color-rgb(0,0,0)"},{"offset":0,"length":111,"style":"bgcolor-transparent"},{"offset":0,"length":111,"style":"fontsize-11pt"},{"offset":0,"length":111,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"eq5m9","text":"The molar mass of an oxygen molecule. is the sum of the molar masses of its 2 atoms. So, 2 × 16 g/mol = 32 g/mol. ","type":"unordered-list-item","depth":0,"inlineStyleRanges":[{"offset":0,"length":113,"style":"color-rgb(0,0,0)"},{"offset":0,"length":113,"style":"bgcolor-transparent"},{"offset":0,"length":113,"style":"fontsize-11pt"},{"offset":0,"length":113,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}}],"entityMap":{}}

3. Use Molar Mass to Convert the Given Amount of the Reactants to Moles

{"blocks":[{"key":"beade","text":"If you want to find out the theoretical yield of a chemical reaction, this means that you have in mind a certain amount of reactants that you're wondering about. You'll use the molar mass calculated above to convert the given amount of each reactant to moles.","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":259,"style":"color-rgb(0,0,0)"},{"offset":0,"length":259,"style":"bgcolor-transparent"},{"offset":0,"length":259,"style":"fontsize-11pt"},{"offset":0,"length":259,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"cu79b","text":"Say you're using 10 grams of hydrogen and 48 grams of oxygen. The calculation will go as follows:","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":97,"style":"color-rgb(0,0,0)"},{"offset":0,"length":97,"style":"bgcolor-transparent"},{"offset":0,"length":97,"style":"fontsize-11pt"},{"offset":0,"length":97,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"21m3j","text":"20 g H2 / (2 g/mol) = 10 moles of hydrogen.","type":"unordered-list-item","depth":0,"inlineStyleRanges":[{"offset":0,"length":43,"style":"color-rgb(0,0,0)"},{"offset":0,"length":43,"style":"bgcolor-transparent"},{"offset":0,"length":43,"style":"fontsize-11pt"},{"offset":0,"length":43,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"e90le","text":"80 g O2 / (32 g/mol) = 2.5 moles of oxygen. ","type":"unordered-list-item","depth":0,"inlineStyleRanges":[{"offset":0,"length":43,"style":"color-rgb(0,0,0)"},{"offset":0,"length":43,"style":"bgcolor-transparent"},{"offset":0,"length":43,"style":"fontsize-11pt"},{"offset":0,"length":43,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}}],"entityMap":{}}

4. Calculate the Molar Ratio between the Reactants

{"blocks":[{"key":"88m8p","text":"The goal of this step is to know the number of molecules of each reactant you're starting with. To determine the ratio between the two reactants, divide the number of moles of the first by the number of moles of the second.","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":223,"style":"color-rgb(0,0,0)"},{"offset":0,"length":223,"style":"bgcolor-transparent"},{"offset":0,"length":223,"style":"fontsize-11pt"},{"offset":0,"length":223,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"cllaf","text":"So, in this example, the molar ratio between hydrogen and oxygen is 5/2.5 = 4. This means that the reaction is starting with 4 times as many hydrogen molecules as oxygen molecules.","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":180,"style":"color-rgb(0,0,0)"},{"offset":0,"length":180,"style":"bgcolor-transparent"},{"offset":0,"length":180,"style":"fontsize-11pt"},{"offset":0,"length":180,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}}],"entityMap":{}}

5. Find the Reaction's Ideal Ratio

{"blocks":[{"key":"ecdm","text":"You can easily determine this by observing the balanced equation of the chemical reaction. Note the coefficient in front of each reactant to find the ratio of molecules required to complete the reaction.","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":203,"style":"color-rgb(0,0,0)"},{"offset":0,"length":203,"style":"bgcolor-transparent"},{"offset":0,"length":203,"style":"fontsize-11pt"},{"offset":0,"length":203,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"9os2p","text":"In the case of the equation 2 H2(g) + O2(g) → 2 H2O(l), you can see that for every oxygen molecule, you need 2 hydrogen molecules.","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":130,"style":"color-rgb(0,0,0)"},{"offset":0,"length":130,"style":"bgcolor-transparent"},{"offset":0,"length":130,"style":"fontsize-11pt"},{"offset":0,"length":130,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}}],"entityMap":{}}

6. Pinpoint the Limiting Reactant

{"blocks":[{"key":"9g45k","text":"The limiting reactant is the one that gets used up first. By identifying the limiting reactant of a chemical reaction, you can determine how long it can occur, and consequently, the theoretical yield to expect.","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":210,"style":"color-rgb(0,0,0)"},{"offset":0,"length":210,"style":"bgcolor-transparent"},{"offset":0,"length":210,"style":"fontsize-11pt"},{"offset":0,"length":210,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"70hf7","text":"In the water formation example, you're starting with 4 times as much hydrogen as oxygen in terms of moles. Since the ideal ratio is 2 times as much hydrogen as oxygen, this means you have more hydrogen than needed. As such, the limiting reactant is oxygen.","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":256,"style":"color-rgb(0,0,0)"},{"offset":0,"length":256,"style":"bgcolor-transparent"},{"offset":0,"length":256,"style":"fontsize-11pt"},{"offset":0,"length":256,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}}],"entityMap":{}}

7. Choose the Desired Product and Determine its Ratio to the Limiting Reactant

{"blocks":[{"key":"a0o6v","text":"If your chemical reaction has more than one product, then you'll have to choose which one you want to work with. In the case of our example, there's only one product.","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":166,"style":"color-rgb(0,0,0)"},{"offset":0,"length":166,"style":"bgcolor-transparent"},{"offset":0,"length":166,"style":"fontsize-11pt"},{"offset":0,"length":166,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"82l84","text":"Once you've selected a product, review the balanced equation and divide the number of molecules of this product by the number of molecules of the limiting reaction.","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":164,"style":"color-rgb(0,0,0)"},{"offset":0,"length":164,"style":"bgcolor-transparent"},{"offset":0,"length":164,"style":"fontsize-11pt"},{"offset":0,"length":164,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"77epu","text":"For 2 H2(g) + O2(g) → 2 H2O(l), the ratio of water to oxygen is 2/1 = 2. ","type":"unordered-list-item","depth":0,"inlineStyleRanges":[{"offset":0,"length":72,"style":"color-rgb(0,0,0)"},{"offset":0,"length":72,"style":"bgcolor-transparent"},{"offset":0,"length":72,"style":"fontsize-11pt"},{"offset":0,"length":72,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}}],"entityMap":{}}

8. Multiply the Ratio by the number of Moles of the Limiting Reactant

{"blocks":[{"key":"cauts","text":"When you do this, you'll get the theoretical yield in moles. So in today's example, the theoretical yield of water is 2 (ratio of water to oxygen) × 2.5 (moles of oxygen) = 5 moles of water.","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":190,"style":"color-rgb(0,0,0)"},{"offset":0,"length":190,"style":"bgcolor-transparent"},{"offset":0,"length":190,"style":"fontsize-11pt"},{"offset":0,"length":190,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}}],"entityMap":{}}

9. Use the Molar Mass of the Product to Convert the Result to Grams

{"blocks":[{"key":"aieec","text":"To determine the theoretical yield in grams, multiply the resulting number of moles by the molar mass of the product. In the water formation example:","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":149,"style":"color-rgb(0,0,0)"},{"offset":0,"length":149,"style":"bgcolor-transparent"},{"offset":0,"length":149,"style":"fontsize-11pt"},{"offset":0,"length":149,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"296k2","text":"The molar mass of a single H2O molecule is 18 grams.","type":"unordered-list-item","depth":0,"inlineStyleRanges":[{"offset":0,"length":52,"style":"color-rgb(0,0,0)"},{"offset":0,"length":52,"style":"bgcolor-transparent"},{"offset":0,"length":52,"style":"fontsize-11pt"},{"offset":0,"length":52,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"d8vi9","text":"5 (number of moles) × 18 = 90 grams.","type":"unordered-list-item","depth":0,"inlineStyleRanges":[{"offset":0,"length":36,"style":"color-rgb(0,0,0)"},{"offset":0,"length":36,"style":"bgcolor-transparent"},{"offset":0,"length":36,"style":"fontsize-11pt"},{"offset":0,"length":36,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}},{"key":"e7u56","text":"As such, 80 grams of oxygen gas with excess hydrogen will theoretically yield 90 grams of water.","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":96,"style":"color-rgb(0,0,0)"},{"offset":0,"length":96,"style":"bgcolor-transparent"},{"offset":0,"length":96,"style":"fontsize-11pt"},{"offset":0,"length":96,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}}],"entityMap":{}}

Conclusion

{"blocks":[{"key":"dfb1","text":"Be sure to try these steps with different equations to get used to the calculations. Also, for chemical reactions with multiple products, you can simply repeat the steps for other products one at a time.","type":"unstyled","depth":0,"inlineStyleRanges":[{"offset":0,"length":203,"style":"color-rgb(0,0,0)"},{"offset":0,"length":203,"style":"bgcolor-transparent"},{"offset":0,"length":203,"style":"fontsize-11pt"},{"offset":0,"length":203,"style":"fontfamily-Arial"}],"entityRanges":[],"data":{}}],"entityMap":{}}