I got 94.07556465 rounded to 94.076% Is this correct?

Ya! You can avoid that much error. But, try to go as far as you can while you are rounding up or calculating. At least to a nearest 100th place.

The article says "Atomic mass is commonly referred to as AMU (atomic mass number)" but atomic mass number is A, not AMU. AMU is atomic mass unit which is 1.66054e-24 g.

Yes, you're right. But carbon is standardized where 12 amu = 12 grams. For our purposes of college courses chem 1 and 2 and standardized tests like the MCAT, we can assume 1 amu = 1g just like we can assume gravity, g = 10 m/s^2.

Main content

Course: Class 11 Chemistry (India) > Unit 1

Lesson 6: Stoichiometry

Stoichiometry article

Google Classroom

In our day to day life, it’s child’s play for us to ‘count’ things or ‘weigh’ things. For example, you can easily count and tell me that there are 4 bananas or there are 6 apples kept on the table. Shopkeepers can easily weigh things on a weighing balance and tell us ‘this packet of rice weighs 2 pounds’ or ‘this packet of tea weighs 500 g’. But things change when you start dealing with atoms and molecules. You cannot simply count or weigh atoms or molecules, as you can do with the common items that you encounter in your daily life. Here it gets a little tricky and something termed as ‘Stoichiometry’ kicks in.

How do you define stoichiometry?

Stoichiometry is the branch of chemistry that deals with the relationship between the relative quantities of substances taking part in a chemical reaction

Let’s write a general chemical reaction where there are two reactants, A and B that react together to form two products, C and D respectively

A + B → C + D

The concept of stoichiometry will help us answer questions like

if ‘x’ grams of A is present, then how many grams of C or D or both will be produced
if ‘y’ grams of B is present, then how many grams of C or D or both will be produced
if ‘x’ grams of A is present and ‘y’ grams of B is present, which is the limiting reactant and which is the reactant present in excess
if ‘x’ grams of A react with ‘y’ grams of B, how much of B will remain unreacted (or unconsumed) after the reaction is over
how many grams of A will produce ‘z’ grams of C

The concept of stoichiometry can be best understood by solving real problems involving chemical reactions, reactants and products

Just two things to remember before we get started:

the first step in any stoichiometric problem is to always ensure that the chemical reaction you are dealing with is balanced,
clarity of the concept of a ‘mole’ and the relationship between ‘amount (grams)’ and ‘moles’.

‘Mole concept’ in brief

Let’s write a simple balanced chemical reaction, where hydrogen gas reacts with oxygen gas to produce water.

2

H

_{2}

O

_{2}

→

2

H

_{2}

O

In the above reaction the prefixes (shown in red) refer to the moles of the molecules involved in the chemical reaction. These are also referred to as coefficients. Here, two moles of hydrogen gas are reacting with one mole of oxygen gas to produce two moles of water. ‘Mole’ simply refers to an enormously big number which is 6.023 X 10

^{23}

Let’s see how a particular element is represented in a periodic table. Let’s pick chlorine (Cl).

As you can see, there are three pieces of information provided here.

Atomic symbol of chlorine
Atomic number of chlorine; this corresponds to the number of protons or electrons present in an atom of chlorine
Atomic mass of chlorine; this value represents the total mass of both the protons and neutrons present in an atom of chlorine. Atomic mass is commonly referred to as A, or atomic mass number.

Rule of thumb is: atomic mass of an element (in grams) = mass of 1 mole of atoms of that element. So in the case of chlorine (Cl), one mole of chlorine atoms (i.e. 6.023 X 10

^{23}

atoms of chlorine) will weigh 35.453 g.

The same principle can be applied to a molecule also, example water (H

_{2}

O).

First we need to calculate the molar mass of H

_{2}

O = 2 (atomic mass of H) + 1(atomic mass of O) = 2(1.008) + 1(15.999) = 18.015 g

So, one mole of water molecules will weigh 18.015 g

The conversion between moles and mass (in grams) should be very clear to us before we move on because this is going to be the basis of all stoichiometric calculations.

Let’s find out how many moles of glucose are present in 25 g of glucose

The first step, as always, is to calculate the molar mass of glucose (C

_{6}

_{12}

_{6}

) = 6 (atomic mass of C) + 12 (atomic mass of H) + 6 (atomic mass of O) = 6 (12.011) + 12 (1.008) + 6 (15.999) = 72.066 + 12.096 + 95.994 = 180.156 grams

So, 1 mole of glucose weighs 180.156 g

i.e. 180.156 g of glucose = 1 mole of glucose molecules

therefore, 1 g of glucose = (1 mole/180.156 g) x 1 g

so, 25 g of glucose = (1 mole/180.156 g) x 25 g = 0.139 moles

Now let’s practice some stoichiometric problems

Problem 1

If 138.6 g of KClO

_{3}

is heated and decomposes completely, how many grams of oxygen gas would be produced?

$K C l O_{3} \to K C l + O_{2}$ ‍

Step 1: Ensure that the reaction is balanced

In this case the reaction is not balanced. After balancing, the equation looks like

$2 K C l O_{3} \to 2 K C l + 3 O_{2}$ ‍

Step 2: Convert grams of KClO $_{3}$ ‍ to moles

For this we first need to calculate the molar mass of KClO

_{3}

= 1 (atomic mass of K) + 1 (atomic mass of Cl) + 3 (atomic mass of O) = 1 (39.098) + 1 (35.453) + 3 (15.999) = 39.098 + 35.453 + 47.997 = 122.548 g

i.e. 122.548 g of KClO

_{3}

= 1 mole of KClO

_{3}

therefore, 1 g of KClO

_{3}

= (1 mole/ 122.548 g) x 1 g

so, 138.6 g of KClO

_{3}

= (1 mole/ 122.548 g) x 138.6 g = 1.130 moles

Step 3: Use mole ratios (from the balanced chemical reaction) to calculate the moles of the relevant reactant or product

In this example, we need to calculate the mass of O

_{2}

produced. So we first need to calculate the moles of O

_{2}

that would be produced from 1.130 moles of KClO

_{3}

Looking at the balanced chemical reaction, we can infer that 2 moles of KClO

_{3}

produce 3 moles of O

_{2}

So, 1.130 moles of KClO

_{3}

will produce [(3 moles of O

_{2}

/ 2 moles of KClO

_{3}

) x 1.130 moles of KClO

_{3}

] = 1.695 moles of O

_{2}

Step 4: Convert moles back to ‘grams’ using molar mass of O $_{2}$ ‍

Molar mass of O

_{2}

= 2 (atomic mass of O) = 2 (15.999) = 31.998 g i.e. 1 mole of O

_{2}

= 31.998 g of O

_{2}

therefore, 1.695 moles of O

_{2}

= 31.998 g/mol x 1.695 mol = 54.236 g

Answer: If 138.6 g of KClO $_{3}$ ‍ is heated and decomposes completely, 54.236 g of oxygen gas will be produced

Problem 2

How many grams of iodine must react to give 4.63 grams of ferric iodide?

$2 F e + 3 I_{2} \to 2 F e I_{3}$ ‍

Step 1: Ensure that the reaction is balanced

The above reaction is balanced!!!

Step 2: Convert grams of FeI $_{3}$ ‍ to moles

For this we first need to calculate the molar mass of FeI

_{3}

= 1 (atomic mass of Fe) + 3 (atomic mass of I) = 1 (55.845) + 3 (126.904) = 55.845 + 380.712 = 436.557 g

i.e. 436.557 g of FeI

_{3}

= 1 mole of FeI

_{3}

therefore, 1 g of FeI

_{3}

= (1 mole/ 436.557 g) x 1 g

so, 4.63 g of FeI

_{3}

= (1 mole/ 436.557 g) x 4.63 g = 0.0106 moles

Step 3: Use mole ratios (from the balanced chemical reaction) to calculate the moles of the relevant reactant or product

In this example, we need to calculate the mass of I

_{2}

that must react. So we first need to calculate the moles of I

_{2}

that would be required to produce 0.0106 moles of FeI

_{3}

Looking at the balanced chemical reaction, we can infer that 3 moles of I

_{2}

are required to produce 2 moles of FeI

_{3}

So, 0.0106 moles of FeI

_{3}

will require [(3 moles of I

_{2}

/ 2 moles of FeI

_{3}

) x 0.0106 moles of FeI

_{3}

] = 0.0159 moles of I

_{2}

Step 4: Convert moles back to ‘grams’ using molar mass of I $_{2}$ ‍

Molar mass of I

_{2}

= 2 (atomic mass of I) = 2 (126.904) = 253.808 g

i.e. 1 mole of I

_{2}

= 253.808 g of I

_{2}

therefore, 0.0159 moles of I

_{2}

= (253.808 x 0.0159) g = 4.035 g

Answer: 4.035 grams of iodine must react to yield 4.63 grams of ferric iodide.

Problem 3

How many grams of H

_{2}

O will be produced when you burn 25 grams of methane?

$C H_{4} + O_{2} \to C O_{2} + H_{2} O$ ‍

Step 1: Ensure that the reaction is balanced

In this case the reaction is not balanced. After balancing, the equation looks like

$C H_{4} + 2 O_{2} \to C O_{2} + 2 H_{2} O$ ‍

Step 2: Convert grams of CH $_{4}$ ‍ to moles

For this we first need to calculate the molar mass of CH

_{4}

= 1 (atomic mass of C) + 4 (atomic mass of H) = 1 (12.011) + 4 (1.008) = 12.011 + 4.032 = 16.043 g

i.e. 16.043 g of CH

_{4}

= 1 mole of CH

_{4}

therefore, 1 g of CH

_{4}

= (1 mole/ 16.043 g) x 1 g

so, 25 g of CH

_{4}

= (1 mole/ 16.043 g) x 25 g = 1.558 moles

Step 3: Use mole ratios (from balanced chemical reaction) to calculate the moles of the relevant reactant or product

In this example, we need to calculate the mass of H

_{2}

O produced. So we first need to calculate the moles of H

_{2}

O that would be produced from 1.558 moles of CH

_{4}

Looking at the balanced chemical reaction, we can infer that 1 mole of CH

_{4}

produces 2 moles of H

_{2}

So, 1.558 moles of CH

_{4}

will produce [(2 moles of H

_{2}

O / 1 mole of CH

_{4}

) x 1.558 moles of CH

_{4}

] = 3.116 moles of H

_{2}

Step 4: Convert moles back to ‘grams’ using molar mass of H $_{2}$ ‍O

Molar mass of H

_{2}

O = 2 (atomic mass of H) + 1 (atomic mass of O) = 2 (1.008) + 1 (15.999) = 18.015 g

i.e. 1 mole of H

_{2}

O = 18.015 g of H

_{2}

therefore, 3.116 moles of H

_{2}

O = (18.015 x 3.116) g = 56.134 g

Answer: 56.134 g of H $_{2}$ ‍O will be produced when we burn 25 grams of methane

After going through the above problems, we should now have a fair idea of how to calculate the amount of a reactant used or the amount of a product formed in a chemical reaction from a given amount of any one of the reactants or products.

Concept of ‘limiting reactant’ and ‘excess reactant’

Limiting Reactant is defined as the reactant in a chemical reaction that limits the amount of product that is formed. The reaction will stop when all of the limiting reactant is consumed.

Excess Reactant is defined as the reactant in a chemical reaction that is in excess and remains unconsumed when a reaction stops because the limiting reactant has been completely consumed.

Let’s understand this concept with an example. Suppose there are 4 kids and 9 pairs of flip-flops. Each kid will wear a pair of flip-flops, leaving behind 5 pairs of flip-flops.

No matter how many pairs of flip-flops there are, only 4 pairs of flip-flops can be worn because there are only 4 kids. Likewise in chemical reaction, when the limiting reagent (kid) is used up, the reaction will stop and the unconsumed excess reagent (pairs of flip-flops) will remain in the reaction mixture.

Now the key task is to be able to identify which reactant is the limiting reactant in a given chemical reaction.

The easiest way to identify which reactant is the ‘limiting reactant’ and which one is the ‘excess reagent’ is by taking the number of moles of each reactant and dividing it by its molar coefficient, as indicated in the balanced chemical reaction. The substance that has the lowest value is the limiting reagent.

As illustrated above, in an ideal world 36 popcorns (product) should pop out of 36 seeds (reactant). In this case, ‘36’ is referred to as the ‘theoretical yield’. But actually, only 27 popcorns pop. Here, ‘27’ is referred to as the ‘actual yield. The percent yield of this popcorn making process would be defined as (

\frac{a c t u a l y i e l d}{t h e o r e t i c a l y i e l d}

x 100) = 27/ 36 x 100 = 75%. In other words, the popcorn making process was only 75% efficient.

The above analogy can be applied to any chemical reaction. As we have learned in the first part of this tutorial, the amount of product formed can be theoretically calculated from the given amount of the limiting reagent and the mole ratios of the reactants and products. This is called the theoretical yield of a reaction.

However, the amount of product actually formed in a reaction is usually less than the theoretical yield and is referred to as the actual yield. This is because reactions often have "side reactions" that also compete for reactants and produce undesired products. To evaluate the efficiency of the reaction, chemists compare the actual and theoretical yields by calculating the percent yield of a reaction:

Percent yield of a reaction =

\frac{a c t u a l y i e l d}{t h e o r e t i c a l y i e l d}

x 100

Let’s attempt a problem now.

Problem 4

Let’s take the example of preparation of nitrobenzene (C

_{6}

_{5}

_{2}

), starting with 35.15 g of benzene (C

_{6}

_{6}

) and an excess of nitric acid (HNO

_{3}

). The balanced chemical reaction for this is

C_{6} H_{6} + H N O_{3} \to C_{6} H_{5} N O_{2} + H_{2} O

Calculate the percent yield if your preparation yielded 52.145g of nitrobenzene.

Let’s first calculate the theoretical yield of C

_{6}

_{5}

_{2}

Step 1: Convert grams of benzene to moles

Molar mass of C

_{6}

_{6}

= 6 (atomic mass of C) + 6 (atomic mass of H) = 6 (12.011) + 6 (1.008) = 72.066 + 6.048 = 78.114 g

i.e. 78.114 g of C

_{6}

_{6}

= 1 mole of C

_{6}

_{6}

therefore, 1 g of C

_{6}

_{6}

= (

\frac{mole}{78.114 g}

) x 1 g

so, 35.15 g of C

_{6}

_{6}

= (

\frac{1 mole}{78.114 g}

) x 35.15 g = 0.449 moles

Step 2: Determine the theoretical yield of nitrobenzene (convert moles of benzene to moles of nitrobenzene)

Looking at the balanced chemical reaction, we can infer that 1 mole of C

_{6}

_{6}

produces 1 mole C

_{6}

_{5}

_{2}

So, 0.449 moles of C

_{6}

_{6}

will theoretically produce

\frac{1 mole of C_{6} H_{5} N O_{2}}{1 mole of C_{6} H_{6}}

= 0.449 moles of C

_{6}

_{5}

_{2}

Step 3: Convert moles back to grams using molar mass of C $_{6}$ ‍H $_{5}$ ‍NO $_{2}$ ‍

Molar mass of C

_{6}

_{5}

_{2}

= 6 (atomic mass of C) + 5 (atomic mass of H) + 1 (atomic mass of N) + 2 (atomic mass of O) = 6 (12.011) + 5 (1.008) + 1 (14.0067) + 2 (15.999) = 72.066 + 5.04 + 14.0067 + 31.998 = 123.111g

i.e. 1 mole of C

_{6}

_{5}

_{2}

= 123.111 grams of C

_{6}

_{5}

_{2}

Therefore, 0.449 moles of C

_{6}

_{5}

_{2}

= (123.111 x 0.449) g = 55.277g

Step 4: Determine the percent yield.

Percent yield of a reaction = actual yield/theoretical yield x 100%

Answer: Percent yield of nitrobenzene = (52.145g/55.277g) x 100% = 94.334%

Want to join the conversation?

Sort by:

Siddarth T
Posted 3 years ago. Direct link to Siddarth T's post “I got 94.07556465 rounded...”
I got 94.07556465 rounded to 94.076%

Is this correct?
Button navigates to signup pageButton navigates to signup page
(5 votes)
Answer
- yash.patel3
  Posted a year ago. Direct link to yash.patel3's post “Ya! You can avoid that mu...”
  Ya! You can avoid that much error. But, try to go as far as you can while you are rounding up or calculating. At least to a nearest 100th place.
  Button navigates to signup page
  (2 votes)
ff142
Posted 8 years ago. Direct link to ff142's post “The article says "Atomic ...”
The article says "Atomic mass is commonly referred to as AMU (atomic mass number)" but atomic mass number is A, not AMU. AMU is atomic mass unit which is 1.66054e-24 g.
Button navigates to signup pageButton navigates to signup page
(2 votes)
Answer
- Isaac Deatherage
  Posted 7 years ago. Direct link to Isaac Deatherage's post “Yes, you're right. But ca...”
  Yes, you're right. But carbon is standardized where 12 amu = 12 grams. For our purposes of college courses chem 1 and 2 and standardized tests like the MCAT, we can assume 1 amu = 1g just like we can assume gravity, g = 10 m/s^2.
  Button navigates to signup page
  (3 votes)
J Maxwell
Posted 7 years ago. Direct link to J Maxwell's post “It skips from Problem 3 t...”
It skips from Problem 3 to Problem 5 and is missing a good chunk of Problem 5.....otherwise some good quality practice.
Button navigates to signup pageButton navigates to signup page
(2 votes)
Answer
Amrit Kaur
Posted 9 years ago. Direct link to Amrit Kaur's post “"therefore, 0.0159 moles ...”
"therefore, 0.0159 moles of I_2
2
start subscript, 2, end subscript = (253.808 x 0.0159) g = 4.035 g
Answer: 3.807 grams of iodine must react to give 4.035 grams of ferric iodide"

Problem 2 the answer should be 4.035g
Button navigates to signup pageButton navigates to signup page
(2 votes)
Answer
Ginny Redgraves
Posted 7 years ago. Direct link to Ginny Redgraves's post “Hi! I'm wondering if you ...”
Hi! I'm wondering if you can show how you round your figures at each step in the calculation. For example, for Problem 2, I got 4.0355472 after I converted moles to grams. If I were to round the answer to the 3 decimal place, I'd get 4.0356 g and not 4.035 g. This makes me very confused, because I'm now wondering if I rounded my answer wrongly. At least, if you indicated how you rounded your answer, I have something to refer to while redoing my calculations.
Button navigates to signup pageButton navigates to signup page
(1 vote)
Answer
Faiza
Posted 5 years ago. Direct link to Faiza's post “Isn't 35.453 the atomic W...”
Isn't 35.453 the atomic WEIGHT of chlorine rather than atomic MASS? I have also seen these terms being used interchangeably. What's the correct term to use in this case?
Button navigates to signup pageButton navigates to signup page
(1 vote)
Answer
LilyReb
Posted 4 years ago. Direct link to LilyReb's post “In problem 1, step 2, how...”
In problem 1, step 2, how did you jump from 122.548g to 138.6g of KClO3?
Button navigates to signup pageComment on LilyReb's post “In problem 1, step 2, how...”
(1 vote)
Answer