assignment stoichiometry a new hope

What is Stoichiometry? Examples and Practice

• The Albert Team
• Last Updated On: March 17, 2024

Have you ever wondered how chemists know exactly how much of each substance to use in a reaction? The answer lies in a fundamental concept called stoichiometry. This crucial aspect of chemistry helps scientists and students alike understand the quantitative relationships in chemical reactions, ensuring that no atom is wasted. In this post, we’ll delve into what stoichiometry is, explore some engaging stoichiometry examples, and provide you with practical stoichiometry practice problems. Whether you’re just getting started or looking to sharpen your skills, this guide will equip you with the knowledge and tools to tackle stoichiometry problems with confidence.

What We Review

What is Stoichiometry?

Stoichiometry is like the recipe for a cake, but instead of baking, we’re dealing with chemical reactions. At its core, stoichiometry is the study of the quantitative relationships between the reactants and products in chemical reactions. When chemists conduct experiments, they need to know how much of each reactant to use and what amount of product to expect. Stoichiometry provides these answers, ensuring that chemical reactions are efficient and effective.

Understanding the Basics

To grasp stoichiometry, you first need to be familiar with a few key concepts:

• Moles: Just like a dozen represents 12 items, a mole is a specific number of molecules (approximately 6.022\times10^{23} molecules, known as Avogadro’s number).
• Molar Mass: This is the weight of one mole of a substance, typically expressed in grams per mole (g/mol). It tells you how much one mole of a substance weighs.
• Mole Ratios: These ratios come from the coefficients of a balanced chemical equation. They tell you the proportion of reactants and products in a reaction.

Real-World Applications

Stoichiometry isn’t just a topic in your chemistry textbook; it has real-world applications. Pharmacists use stoichiometry to mix medications, environmental scientists use it to track pollutant levels, and engineers apply it to design reactors. Understanding stoichiometry means you’re learning a skill that scientists and professionals use daily to make the world work.

Stoichiometry Examples

Understanding stoichiometry becomes clearer with practical examples. Let’s explore a couple of scenarios to illustrate how stoichiometry helps us solve chemical equations.

Example 1: Water Formation

Consider the chemical reaction where hydrogen and oxygen combine to form water:

If we start with 36.0 grams of H_2 and have an excess of O_2 , how many grams of H_2O can we produce? The molar mass of H_2 is 2.02 g/mol, and for H_2O , it’s 18.02 g/mol.

First, convert the mass of H_2 to moles:

Applying the mole ratio from the balanced equation, calculate the moles of H_2O produced:

Then, convert moles of H_2O to grams:

Thus, 36.0 grams of H_2 can produce 321.15 grams of H_2O .

Example 2: Ammonia Synthesis

Now, let’s examine another one of our stoichiometry examples. This problem examines the synthesis of ammonia ( NH_3 ) from nitrogen ( N_2 ) and hydrogen ( H_2 ) in a reaction conducted at STP (standard temperature and pressure):

If we begin with 22.4 liters of N_2 (1 mole at STP) and have an excess of H_2 , how many grams of NH_3 can be produced? The molar mass of NH_3 is 17.03 g/mol.

Using the mole ratio from the chemical equation:

Then, convert moles of NH_3 to grams:

This result shows that 22.4 liters of N_2 at STP can produce 34.06 grams of NH_3 .

Guided Stoichiometry Practice

To master stoichiometry, practice is key. In this section, we’ll explore strategies to approach stoichiometry problems and provide some guided practice examples. These strategies will help you develop a systematic approach to solving stoichiometry questions, enhancing your problem-solving skills.

Understanding the Approach

Identify What’s Given and What’s Unknown: Start by determining what information is provided in the problem and what you need to find. This could be the amount of reactants, products, or even the coefficients in a balanced chemical equation.

Write the Balanced Chemical Equation: Ensure you have the correct balanced chemical equation for the reaction you’re analyzing. This step is crucial as it sets the foundation for the stoichiometric calculations.

Use Mole Ratios: Utilize the coefficients from the balanced equation to establish mole ratios. These ratios are the heart of stoichiometry, allowing you to convert between amounts of reactants and products.

Perform the Calculations: With the mole ratios, you can perform calculations to find the unknown quantities. Remember to keep track of your units, ensuring they are consistent throughout the calculation.

Guided Example

Let’s practice with a guided example. Consider the combustion of propane ( C_3H_8 ) in oxygen to produce carbon dioxide and water:

Suppose you start with 44.1 grams of propane ( C_3H_8 ) and have an excess of oxygen ( O_2 ). How many grams of carbon dioxide ( CO_2 ) are produced?

Given: 44.1 grams of C_3H_8 The molar mass of C_3H_8 is 44.1 g/mol, and for CO_2 , it’s 44.0 g/mol.

Unknown: Grams of CO_2 produced

First, convert the mass of C_3H_8 to moles:

Mole Ratio: From the balanced equation, the mole ratio of C_3H_8 to CO_2 is 1:3.

Calculation:

Now, convert the moles of CO_2 to grams:

So, 44.1 grams of C_3H_8 produces 132.0 grams of CO_2 .

Stoichiometry Practice Problems

Now that you’ve seen stoichiometry in action through examples, it’s time to test your knowledge with some practice problems. Try to solve these problems on your own before checking the solutions provided. This will help you understand stoichiometry better and prepare you for similar problems in the future.

When magnesium burns in the air, it reacts with oxygen to form magnesium oxide. The balanced chemical equation for this reaction is:

If you start with 24.0 grams of magnesium ( Mg ) and have an excess of oxygen ( O_2 ), how many grams of magnesium oxide ( MgO ) will be produced? The molar mass of Mg is 24.3 g/mol, and MgO is 40.3 g/mol.

Hydrochloric acid reacts with sodium hydroxide to produce sodium chloride and water in the following reaction:

If 36.5 grams of hydrochloric acid ( HCl ) react with sufficient sodium hydroxide ( NaOH ), how many grams of sodium chloride ( NaCl ) are produced? Assume the molar mass of HCl is 36.5 g/mol and NaCl is 58.4 g/mol.

Nitrogen gas can react with hydrogen gas to form ammonia through the following reaction:

Calculate the amount of ammonia ( NH_3 ) produced (in grams) when 28.0 liters of nitrogen gas ( N_2 ) react with an excess of hydrogen gas ( H_2 ) at STP. The molar mass of NH_3 ​ is [/latex]17.03[/latex] g/mol.

Stoichiometry Practice Problems Solutions

Here, we’ll go through the solutions to the stoichiometry practice problems we presented earlier. These solutions will help you understand the step-by-step process involved in solving stoichiometry problems.

Solution to Problem 1:

For the reaction 2Mg + O_2 \rightarrow 2MgO , we’re starting with 24.0 grams of Mg . First, convert the mass of Mg to moles:

Using the stoichiometry of the reaction, convert moles of Mg to moles of MgO (the ratio is 1:1):

Now, convert moles of MgO to grams:

So, 24.0 grams of Mg will produce 39.81 grams of MgO .

Solution to Problem 2:

For the reaction HCl + NaOH \rightarrow NaCl + H_2O , we start with 36.5 grams of HCl . Convert the mass of HCl to moles:

The mole ratio of HCl to NaCl is 1:1, so:

Now, convert moles of NaCl to grams:

Thus, 36.5 grams of HCl will produce 58.44 grams of NaCl .

Solution to Problem 3:

For the reaction N_2 + 3H_2 \rightarrow 2NH_3 , starting with 28.0 liters of N_2 ​ at STP (which corresponds to 1.00 mole of N_2 ​):

Using the stoichiometry of the reaction (1:2 ratio of N_2 to NH_3 ):

Convert moles of NH_3 ​ to grams:

So, 28.0 liters of N_2 ​ will produce 42.575 grams of NH_3 ​.

Stoichiometry is a cornerstone concept in chemistry that enables us to predict and quantify the outcomes of chemical reactions. By understanding the relationships between reactants and products, chemists can conduct reactions efficiently and effectively. Throughout this post, we’ve explored what stoichiometry is, provided clear examples, and offered practice problems to help you build your skills.

Remember, the key to mastering stoichiometry lies in practice. The more you work through problems, the more intuitive these concepts will become. Whether you’re balancing equations, calculating molar masses, or determining the amounts of reactants and products, each step you take strengthens your understanding of chemistry.

We encourage you to revisit these examples and practice problems regularly, and don’t hesitate to explore more complex scenarios as you become more comfortable with the basics. Stoichiometry is not just a topic for the classroom; it’s a tool that scientists use to make real-world decisions in industries, research, and environmental science.

So, keep practicing, stay curious, and remember that every chemical equation tells a story of transformation and interaction. With stoichiometry, you’re equipped to understand and narrate these fascinating stories of science.

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Chemistry Steps

General Chemistry

Stoichiometry.

This is a comprehensive, end-of-chapter set of practice problems on stoichiometry that covers balancing chemical equations, mole-ratio calculations, limiting reactants, and percent yield concepts. 

The links to the corresponding topics are given below.

• The Mole and Molar Mass
• Molar Calculations
• Percent Composition and Empirical Formula
• Stoichiometry of Chemical Reactions

Limiting Reactant

• Reaction/Percent Yield
• Stoichiometry Practice Problems

Balance the following chemical equations:

a) HCl + O 2 → H 2 O + Cl 2

b) Al(NO 3 ) 3 + NaOH → Al(OH) 3 + NaNO 3

c) H 2 + N 2 → NH 3

d) PCl 5 + H 2 O → H 3 PO 4 + HCl

e) Fe + H 2 SO 4 → Fe 2 (SO 4 ) 3 + H 2

f) CaCl 2 + HNO 3 → Ca(NO 3 ) 2 + HCl

g) KO 2 + H 2 O → KOH + O 2 + H 2 O 2

h) Al + H 2 O → Al 2 O 3 + H 2

i) Fe + Br 2 → FeBr 3

j) Cu + HNO 3 → Cu(NO 3 ) 2 + NO 2 + H 2 O

k) Al(OH) 3 → Al 2 O 3 + H 2 O

l) NH 3 + O 2 → NO + H 2 O

m) Ca(AlO 2 ) 2 + HCl → AlCl 3 + CaCl 2 + H 2 O

n) C 5 H 12 + O 2 → CO 2 + H 2 O

o) P 4 O 10 + H 2 O → H 3 PO 4

p) Na 2 CrO 4 + Pb(NO 3 ) 2 → PbCrO 4 + NaNO 3

q) MgCl 2 + AgNO 3 → AgCl + Mg(NO 3 ) 2

r) KClO 3 → KClO 4 + KCl

s) Ca(OH) 2 + H 3 PO 4 → Ca 3 (PO 4 ) 2 + H 2 O

Consider the balanced equation:

C 5 H 12 + 8 O 2 → 5CO 2 + 6H 2 O

Complete the table showing the appropriate number of moles of reactants and products.

How many grams of CO 2  and H 2 O are produced from the combustion of 220. g of propane (C 3 H 8 )?

C 3 H 8 (g) + 5O 2 (g) → 3CO 2 (g) + 4H 2 O(g)

How many grams of CaCl 2 can be produced from 65.0 g of Ca(OH) 2 according to the following reaction,

Ca(OH) 2 + 2HCl → CaCl 2 + 2H 2 O

How many moles of oxygen are formed when 75.0 g of Cu(NO 3 ) 2 decomposes according to the following reaction?

2Cu(NO 3 ) 2   → 2CuO + 4NO 2  + O 2

How many grams of MnCl 2  can be prepared from 52.1 grams of MnO 2 ?

MnO 2  + 4HCl → MnCl 2  + Cl 2  + 2H 2 O

Determine the mass of oxygen that is formed when an 18.3-g sample of potassium chlorate is decomposed according to the following equation:

2KClO 3 (s) → 2KCl(s) + 3O 2 (g).

How many grams of H 2 O will be formed when 48.0 grams H 2 are mixed with excess hydrogen gas?

2H 2  + O 2 → 2H 2 O

Consider the chlorination reaction of methane (CH4):

CH 4 (g) + 4Cl 2 (g) → CCl 4 (g) + 4HCl(g)

How many moles of CH 4 were used in the reaction if 51.9 g of CCl4 were obtained?

How many grams of Ba(NO 3 ) 2 can be produced by reacting 16.5 g of HNO 3 with an excess of Ba(OH) 2 ?

Ethanol can be obtained by fermentation – a complex chemical process breaking down glucose to ethanol and carbon dioxide.

                                                  C 6 H 12 O 6    →    2C 2 H 5 OH   +    2CO 2                                                       glucose                   ethanol

How many mL of ethanol (d =0.789 g/mL) can be obtained by this process starting with 286 g of glucose?

36.0 g of butane (C 4 H 10 ) was burned in an excess of oxygen and the resulting carbon dioxide (CO 2 ) was collected in a sealed vessel.

2C 4 H 10 + 13O 2 → 8CO 2 + 10H 2 O

How many grams of LiOH will be necessary to consume all the CO 2 from the first reaction?

2LiOH + CO 2 → Li 2 CO 3 + H 2 O

13. Which statement about limiting reactant is correct?

a) The limiting reactant is the one in a smaller quantity.

b) The limiting reactant is the one in greater quantity.

c) The limiting reactant is the one producing less product.

d) The limiting reactant is the one producing more product.

Find the limiting reactant for each initial amount of reactants.

4NH 3 + 5O 2 → 4NO + 6H 2 O

a) 2 mol of NH 3 and 2 mol of O 2

b) 2 mol of NH 3 and 3 mol of O 2

c) 3 mol of NH 3 and 3 mol of O 2

d) 3 mol of NH 3 and 2 mol of O 2

Note:  This is not a multiple-choice question. Each row represents a separate question where you need to determine the limiting reactant.

How many g of hydrogen are left over in producing ammonia when 14.0 g of nitrogen is reacted with 8.0 g of hydrogen?

N 2 (g) + 3 H 2 (g) → 2 NH 3 (g)

How many grams of PCl 3 will be produced if 130.5 g Cl 2 is reacted with 56.4 g P 4 according to the following equation?

6Cl 2 (g) + P 4 (s) → 4PCl 3 (l)

How many grams of sulfur can be obtained if 12.6 g H 2 S is reacted with 14.6 g SO 2 according to the following equation?

2H 2 S(g) + SO 2 (g) → 3S(s) + 2H 2 O(g)

The following equation represents the combustion of octane, C 8 H 18 , a component of gasoline:

2C 8 H 18 (g) + 25O 2 (g) → 16CO 2 (g) + 18H 2 O(g)

Will 356 g of oxygen be enough for the complete combustion of 954 g of octane?

When 140.0 g of AgNO 3 was added to an aqueous solution of NaCl, 86.0 g of AgCl was collected as a white precipitate. Which salt was the limiting reactant in this reaction? How many grams of NaCl were present in the solution when AgNO 3 was added?

AgNO 3 (aq) + NaCl(aq) → AgCl(s) + NaNO 3 (aq)

Consider the reaction between MnO 2 and HCl:

MnO 2 + 4HCl → MnCl 2 + Cl 2 + 2H 2 O

What is the theoretical yield of MnCl 2 in grams when 165 g of MnO 2 is added to a solution containing 94.2 g of HCl?

Percent Yield

21. In a chemistry experiment, a student obtained 5.68 g of a product. What is the percent yield of the product if the theoretical yield was 7.12 g?

When 38.45 g CCl 4 is reacted with an excess of HF, 21.3 g CCl 2 F 2 is obtained. Calculate the theoretical and percent yields of this reaction.

CCl 4 + 2HF → CCl 2 F 2 + 2HCl

Iron(III) oxide reacts with carbon monoxide according to the equation:

Fe 2 O 3 ( s ) + 3CO( g ) → 2Fe( s ) + 3CO 2 ( g )

What is the percent yield of this reaction if 623 g of iron oxide produces 341 g of iron?

Determine the percent yield of the reaction if 77.0 g of CO 2  are formed from burning 2.00 moles of C 5 H 12  in 4.00 moles of O 2 .

C 5 H 12 + 8 O 2 → 5CO 2  + 6H 2 O

The percent yield for the following reaction was determined to be 84%:

N 2 ( g ) + 2H 2 ( g ) → N 2 H 4 ( l )

How many grams of hydrazine (N 2 H 4 ) can be produced when 38.36 g of nitrogen reacts with 6.68 g of hydrogen?

Silver metal can be prepared by reducing its nitrate, AgNO 3  with copper according to the following equation:

Cu( s ) + 2AgNO 3 ( aq ) → Cu(NO 3 ) 2 ( aq ) + 2Ag( s )

What is the percent yield of the reaction if 71.5 grams of Ag was obtained from 132.5 grams of AgNO 3  ?

Industrially, nitric acid is produced from ammonia by the Ostwald process in a series of reactions:

4NH 3 ( g ) + 5O 2 ( g ) → 4NO( g ) + 6H 2 O( l )

2NO( g ) + O 2 ( g ) → 2NO 2 ( g )

2NO 2 ( g ) + H 2 O( l ) → HNO 3 ( aq ) + HNO 2 ( aq )

Considering that each reaction has an 85% percent yield, how many grams of NH 3 must be used to produce 25.0 kg of HNO 3 by the above procedure?

Aspirin (acetylsalicylic acid) is widely used to treat pain, fever, and inflammation. It is produced from the reaction of salicylic acid with acetic anhydride. The chemical equation for aspirin synthesis is shown below:

In one container, 10.00 kg of salicylic acid is mixed with 10.00 kg of acetic anhydride.

a)  Which reactant is limiting? Which is in excess? b)  What mass of excess reactant is left over? c)  What mass of aspirin is formed assuming 100% yield (Theoretical yield)? d)  What mass of aspirin is formed if the reaction yield is 70.0% ? e)  If the actual yield of aspirin is 11.2 kg, what is the percent yield? f)  How many kg of salicylic acid is needed to produce 20.0 kg of aspirin if the reaction yield is 85.0% ?

3 thoughts on “Stoichiometry Practice Problems”

You forgot the subscript 3 for O in the molecular formula for acetic anhydride and the reaction is not balanced as written. For part F) it’s 18.1 kg and not1.81 kg as written in the final line of the solution.

Thanks for letting me know! Fixed.

You’re welcome!

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