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Magazine Chapter 17: Problem 54
\(0.6 \mathrm{~g}\) of an organic compound was analysed: The ammonia thus produced was absorbed in \(90 \mathrm{~mL} \mathrm{~N} / 9 \mathrm{H}_{2} \mathrm{SO}_{4}\). The remaining acid required \(20 \mathrm{~mL} 0.1 \mathrm{NNaOH}\) for neutralisation. Calculate the percentage of nitrogen in the compound.
Short Answer
Expert verified
The percentage of nitrogen in the compound is 18.67\%.
Step by step solution
01
Understand the Problem
The goal is to find the percentage of nitrogen in the organic compound. First, ammonia produced by the decomposition of the organic compound is absorbed by a certain amount of sulfuric acid. The remaining sulfuric acid is neutralized using sodium hydroxide. These reactions will help us calculate the amount of nitrogen.
02
Determine Total and Consumed Acid
Calculate the total moles of sulfuric acid initially taken and the moles neutralized by sodium hydroxide. Use the formula for normality and volume: \(\text{moles} = \text{Normality} \times \frac{\text{Volume (in L)}}{2}\) (because sulfuric acid is dibasic).
03
Calculate Bicarbonate Moles Neutralized by NaOH
Neutralization by sodium hydroxide gives:\(\text{Moles neutralized} = 0.1 \times \frac{20}{1000} = 0.002\text{ N}\). Thus,Total moles of \( \text{H}_2\text{SO}_4 \) = \( \frac{90}{1000} \times \frac{1}{9} \text{ N} \)= \(0.01\text{ N}\)Remaining free acid: \((0.01 - 0.002)/2 = 0.004\text{ moles of }\text{H}_2\text{SO}_4\) .
04
Calculate Moles of Ammonia
The moles of sulfuric acid consumed by ammonia equal the moles of ammonium produced: \(0.004 \text{ moles}\). Since 1 mole of \( \text{H}_2\text{SO}_4 \) reacts with 2 moles of ammonia, we calculate \( 0.004 \times 2 = 0.008 \text{ moles of ammonia produced.}\)
05
Calculate Mass of Nitrogen in Ammonia
The molar mass of nitrogen is 14 g/mol. Therefore, the mass of nitrogen in the ammonia produced is \(0.008 \text{ moles of ammonia} \times 14 \text{ g/mol} = 0.112 \text{ g}\).
06
Calculate Percentage of Nitrogen
Calculate the percentage of nitrogen in the original organic compound using the formula:\[ \text{Percentage of N} = \left( \frac{\text{Mass of N}}{\text{Initial mass of compound}} \right) \times 100 \% \] \(= \left( \frac{0.112}{0.6} \right) \times 100 = 18.67 \%\)
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Key Concepts
These are the key concepts you need to understand to accurately answer the question.
Organic Compounds Analysis
Analyzing organic compounds is essential in determining their elemental composition, crucial for understanding the structure and properties of the compound.
One common method of analysis involves breaking down the compound and measuring the amounts of different elements it contains.
In this exercise, we dealt with an organic compound containing nitrogen. By decomposing the compound, we produced ammonia, which was then analyzed to find the nitrogen content.
This approach helps identify how much nitrogen is present relative to the compound鈥檚 total mass, enabling calculation of the percentage of nitrogen.
One common method of analysis involves breaking down the compound and measuring the amounts of different elements it contains.
In this exercise, we dealt with an organic compound containing nitrogen. By decomposing the compound, we produced ammonia, which was then analyzed to find the nitrogen content.
This approach helps identify how much nitrogen is present relative to the compound鈥檚 total mass, enabling calculation of the percentage of nitrogen.
Neutralization Reaction
A neutralization reaction occurs when an acid reacts with a base, producing water and a salt.
In the context of the nitrogen content calculation, neutralization is used to assess the amount of sulfuric acid remaining after reacting with ammonia.
In this problem, sulfuric acid absorbed the ammonia, and any leftover acid was neutralized with sodium hydroxide (NaOH).
The resulting reaction allows for determining how much of the acid was initially unavailable, providing a clue to the amount of ammonia produced by the compound.
In the context of the nitrogen content calculation, neutralization is used to assess the amount of sulfuric acid remaining after reacting with ammonia.
In this problem, sulfuric acid absorbed the ammonia, and any leftover acid was neutralized with sodium hydroxide (NaOH).
The resulting reaction allows for determining how much of the acid was initially unavailable, providing a clue to the amount of ammonia produced by the compound.
Sulfuric Acid Usage
Sulfuric acid is a powerful acid often used to absorb ammonia in nitrogen content analysis.
It reacts with ammonia to form ammonium sulfate, helping to estimate the ammonia generated from the organic compound breakdown.
In this exercise, we start with a known quantity of sulfuric acid. After some of it reacts with ammonia, the remaining amount is determined by titration with NaOH.
This difference reveals the volume of sulfuric acid that interacted directly with the ammonia, enabling further calculations.
It reacts with ammonia to form ammonium sulfate, helping to estimate the ammonia generated from the organic compound breakdown.
In this exercise, we start with a known quantity of sulfuric acid. After some of it reacts with ammonia, the remaining amount is determined by titration with NaOH.
This difference reveals the volume of sulfuric acid that interacted directly with the ammonia, enabling further calculations.
Ammonia Production
Ammonia production is a significant step in determining the nitrogen content in an organic compound.
When the compound decomposes, it liberates nitrogen in the form of ammonia, which then interacts with the sulfuric acid.
In the given exercise, we calculated ammonia moles by looking at how much sulfuric acid it neutralized.
We learned that each mole of sulfuric acid reacted corresponds to two moles of ammonia. Therefore, finding the moles of ammonia gives a clear picture of the nitrogen in the compound.
When the compound decomposes, it liberates nitrogen in the form of ammonia, which then interacts with the sulfuric acid.
In the given exercise, we calculated ammonia moles by looking at how much sulfuric acid it neutralized.
We learned that each mole of sulfuric acid reacted corresponds to two moles of ammonia. Therefore, finding the moles of ammonia gives a clear picture of the nitrogen in the compound.
Percent Composition
The percent composition of a chemical compound refers to the percentage by mass of each element in the compound.
For this exercise, the main focus is on the nitrogen composition.
Once the mass of nitrogen in the compound is obtained, calculating the percentage is straightforward:
We divide the mass of the nitrogen by the total mass of the organic compound and multiply by 100.
The goal is to express how much of the compound is made up of nitrogen, which in this problem is found to be 18.67%.
For this exercise, the main focus is on the nitrogen composition.
Once the mass of nitrogen in the compound is obtained, calculating the percentage is straightforward:
We divide the mass of the nitrogen by the total mass of the organic compound and multiply by 100.
The goal is to express how much of the compound is made up of nitrogen, which in this problem is found to be 18.67%.
Stoichiometry
Stoichiometry is the quantitative relationship between reactants and products in a chemical reaction.
It is fundamental in solving problems involving chemical reactions, like the nitrogen content calculation.
In this analysis, stoichiometry helps relate the moles of sulfuric acid with those of ammonia. Since one mole of sulfuric acid reacts with two moles of ammonia, understanding this ratio is key.
It allows us to compute the precise quantity of nitrogen in the compound by following these balanced reaction pathways.
It is fundamental in solving problems involving chemical reactions, like the nitrogen content calculation.
In this analysis, stoichiometry helps relate the moles of sulfuric acid with those of ammonia. Since one mole of sulfuric acid reacts with two moles of ammonia, understanding this ratio is key.
It allows us to compute the precise quantity of nitrogen in the compound by following these balanced reaction pathways.
