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Stoichiometry 6: Finding the number of moles of water produced when 2 moles of oxygen react with excess hydrogen
Introduction
Here’s a practical thought:
Every time a car engine burns fuel, water is produced as a byproduct. Similarly, many industrial chemical reactions also produce water as a byproduct.
But how do chemists predict exactly how much water will be formed from a given amount of reactant?
This is where stoichiometry becomes more than a classroom exercise — it becomes a vital tool for industry.
Today, we’ll solve a classic problem:
How many grams of water are produced when 10 moles of ammonia react with excess oxygen?
We’ll walk through it slowly, step by step, so you not only get the right answer but also understand why each step matters.
Problem Statement — What Are We Solving?
Let’s start with the question:
How many grams of water (H₂O) are produced when 10 moles of ammonia (NH₃) react with excess oxygen (O₂)?
You’re also given:
- Balanced chemical equation:
4 NH₃ + 5 O₂ → 4 NO + 6 H₂O - Molar mass of water (H₂O): 18 g/mol
The key insight is recognizing that ‘excess oxygen’ implies oxygen will not limit the reaction. As a result, ammonia must be the limiting reactant.
With this in mind, we can now focus exclusively on how ammonia stoichiometrically converts to water.
Step 1: Understand the Balanced Chemical Equation
First, the balanced equation serves as your starting point for every stoichiometry problem. Then, it clearly shows you the exact proportions in which reactants combine and products form.
Here’s what it tells us:
- 4 moles of ammonia react with
- 5 moles of oxygen to produce
- 4 moles of nitrogen monoxide (NO) and
- 6 moles of water (H₂O).
For this problem, the important relationship is between ammonia and water:
4 NH₃ : 6 H₂O
This mole ratio means:
For every 4 moles of ammonia, the reaction produces 6 moles of water.
Step 2: Use the Mole Ratio to Calculate Moles of Water Produced
Now that we have the mole ratio, we can use it to find out how much water is formed when 10 moles of ammonia are used.
Here’s how to set it up:
(6 moles H₂O / 4 moles NH₃) × 10 moles NH₃ = 15 moles H₂O
What does this calculation mean?
For every 4 moles of ammonia, you get 6 moles of water.
So, if you have 10 moles of ammonia, you scale up proportionally and produce 15 moles of water.
This step is where many students make mistakes.
They either use the wrong ratio or mix up the direction of the calculation.
Always start from the balanced equation —
It’s your cheat sheet for getting this right.
Step 3: Convert Moles of Water to Grams
The problem specifically asks for the mass of water in grams, not just moles.
To make this conversion, you need to use the molar mass of water, which is given as 18 g/mol.
The formula is simple:
Mass = Moles × Molar Mass
Substitute the values:
Mass = 15 moles × 18 g/mol = 270 grams
So, the reaction will produce 270 grams of water when 10 moles of ammonia react with excess oxygen.
This is the final numerical answer.
But the method behind it is what’s important to understand.
Step 4: Common Mistakes to Watch Out For
Even though the math is straightforward, students frequently lose marks on these kinds of questions because of simple errors.
Here are the common pitfalls:
- Confusing Which Substance is Being Asked For
Sometimes students accidentally calculate how much oxygen is used instead of water produced.
Always double-check what the question is actually asking. - Using the Wrong Molar Mass
Forgetting that H₂O has a molar mass of 18 g/mol or mistakenly using the molar mass of ammonia can derail your answer. - Ignoring Mole-to-Mole Ratios
Assuming a 1:1 conversion between ammonia and water is a classic mistake.
The balanced equation gives you the correct ratio — always refer back to it. - Forgetting to Convert Moles to Grams
Giving your answer in moles when grams are requested is an oversight that costs easy marks.
As a result of staying mindful of these, you can approach every stoichiometry problem with more accuracy and confidence.
Step 5: Real-World Application — Why Does This Matter?
In fact, in industrial chemistry, large-scale manufacturing processes often produce water as a byproduct.
For example:
- In fertilizer production, ammonia reacts with oxygen in processes like the Ostwald method to produce nitric acid, with water as a byproduct.
- In waste gas treatment plants, engineers monitor water formation to control emissions and prevent system malfunctions.
- In chemical reactors, predicting how much water will form helps maintain proper pressure and temperature controls.
Overproduction of water can lead to condensation issues, clogging systems, or even triggering safety hazards.
On the other hand, underestimating it can result in inefficient processes and wasted resources.
That’s why, accurate stoichiometric calculations are essential — they allow facilities to:
- Design better drainage and purification systems.
- Optimize reactant usage for cost efficiency.
- Predict energy changes since water formation often releases heat.
So while this problem might seem like a simple exercise, industries worldwide actively apply the principles behind it.
Step 6: Final Wrap-Up — What You’ve Learned
Let’s bring everything together.
At this point, you have calculated how many grams of water the reaction between 10 moles of ammonia and excess oxygen produces.
The balanced equation (4 NH₃ + 5 O₂ → 4 NO + 6 H₂O) revealed the crucial mole ratio:
4 NH₃ : 6 H₂O
Using this ratio, you calculated:
(6 ÷ 4) × 10 = 15 moles of water produced.
Finally, you converted moles to grams:
15 moles × 18 g/mol = 270 grams of water.
Furthermore, this problem shows how chemists use balanced equations and mole ratios to predict the mass of products formed in a reaction — a fundamental skill in chemistry
In addition, understanding this method builds a foundation for more complex topics like reaction yields, limiting reactants, and even energy calculations.
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At www.copychemistry.com, we specialize in helping students go beyond memorizing formulas.
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Here’s what you’ll get:
- Interactive lessons that make mole ratios and conversions intuitive.
- Practice problems designed to mimic real exam questions.
- Feedback that explains both your mistakes and how to fix them.
- Support to help you approach tests with confidence — not anxiety.
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CopyChemistry is your go-to resource.
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