Determining Molecular Formula Given Empirical Formula NH₃ And Mr 17

by Scholario Team 68 views

Hey guys! Ever stumbled upon a chemistry problem that looks like a jumbled mess of letters and numbers? Don't worry, we've all been there! Today, we're going to break down a classic problem: figuring out the molecular formula of a compound when you're given its empirical formula and its molar mass (Mr). In this case, we have a compound with the empirical formula NH₃ and an Mr of 17. Sounds intimidating? Trust me, it's not! Let's dive in and make sense of it together.

Understanding Empirical and Molecular Formulas

Before we jump into solving the problem, let's quickly refresh our understanding of what empirical and molecular formulas actually represent. This is super important, so pay close attention!

  • Empirical Formula: Think of the empirical formula as the simplest, most basic recipe for a compound. It tells you the smallest whole-number ratio of atoms of each element in the compound. In our case, NH₃ tells us that for every 1 nitrogen (N) atom, there are 3 hydrogen (H) atoms. It's like saying, “Hey, the basic building block of this compound is one N and three Hs.”
  • Molecular Formula: Now, the molecular formula is the real deal. It tells you the actual number of atoms of each element present in a molecule of the compound. It's the complete recipe, showing you exactly how many of each ingredient you need. The molecular formula is always a whole-number multiple of the empirical formula. This multiple is the key to solving our problem!

So, the empirical formula is the simplified version, and the molecular formula is the full, accurate picture. Make sense? Great! Let's move on.

Steps to Determine the Molecular Formula

Alright, now that we've got the basics down, let's tackle the problem step-by-step. I'll break it down into easy-to-follow instructions, so you can conquer similar problems in the future.

  1. Calculate the Empirical Formula Mass: The first step is to find the mass of the empirical formula unit. This is just like calculating the molar mass, but you're using the empirical formula instead of the molecular formula. We'll use the periodic table to find the atomic masses of each element.

    • For NH₃:
      • Nitrogen (N): Atomic mass ≈ 14 amu (atomic mass units)
      • Hydrogen (H): Atomic mass ≈ 1 amu
    • Empirical formula mass = (1 × N) + (3 × H) = (1 × 14) + (3 × 1) = 14 + 3 = 17 amu

    So, the empirical formula mass of NH₃ is 17 amu. This means one “unit” of NH₃, in its simplest ratio, weighs 17 amu.

  2. Determine the Multiplication Factor: This is the crucial step! We need to find out how many times the empirical formula “fits” into the molecular formula. To do this, we'll divide the molar mass (Mr) of the compound by the empirical formula mass.

    • Multiplication factor = Mr / Empirical formula mass = 17 / 17 = 1

    Aha! The multiplication factor is 1. This tells us that the molecular formula is simply one times the empirical formula. This is a key step, guys, so make sure you understand the logic behind it.

  3. Calculate the Molecular Formula: Now, the final step! To find the molecular formula, we multiply the subscripts in the empirical formula by the multiplication factor we just calculated.

    • Molecular formula = (NH₃) × 1 = NH₃

    In this case, since the multiplication factor is 1, the molecular formula is the same as the empirical formula. That's it!

Solution and Explanation

Therefore, the molecular formula of the compound is NH₃. This means that the actual molecule consists of one nitrogen atom and three hydrogen atoms. In this specific instance, the empirical and molecular formulas are the same because the simplest ratio of atoms is also the actual ratio in the molecule. Isn't that neat?

Let's recap why this happened. The molar mass (Mr) of the compound was equal to the empirical formula mass. This indicates that the compound exists as the simplest possible unit represented by the empirical formula. If the Mr were, say, 34, then the multiplication factor would have been 2 (34 / 17 = 2), and the molecular formula would be N₂H₆.

Key Takeaways and Tips

Before we wrap up, let's highlight some key takeaways and tips that will help you tackle similar problems with confidence.

  • Understanding the Definitions: Make sure you have a solid grasp of the difference between empirical and molecular formulas. This is the foundation for solving these types of problems. Remember, empirical is the simplest ratio, and molecular is the actual number of atoms.
  • Step-by-Step Approach: Break down the problem into smaller, manageable steps. Calculate the empirical formula mass, find the multiplication factor, and then determine the molecular formula. This systematic approach will prevent you from getting lost in the process.
  • Periodic Table is Your Friend: The periodic table is your best buddy in chemistry! Use it to find the atomic masses of elements. These values are crucial for calculating empirical and molecular masses.
  • Check Your Work: Always double-check your calculations to avoid silly mistakes. A small error in the beginning can lead to a wrong answer at the end.
  • Practice Makes Perfect: The more you practice, the better you'll become at solving these problems. Try working through different examples with varying empirical formulas and molar masses.

Practice Problems

To solidify your understanding, here are a couple of practice problems for you to try:

  1. A compound has an empirical formula of CH₂ and an Mr of 42. Determine its molecular formula.
  2. A compound has an empirical formula of HO and an Mr of 34. Determine its molecular formula.

Try solving these on your own, and feel free to share your answers in the comments below! Let's help each other learn.

Conclusion

So, there you have it! We've successfully determined the molecular formula of a compound given its empirical formula and molar mass. It might have seemed tricky at first, but by breaking it down into steps and understanding the underlying concepts, it becomes much more manageable. Chemistry can be like solving a puzzle, and it's so rewarding when you crack the code!

Remember, the key is to understand the definitions, follow the steps, and practice, practice, practice! You've got this, guys! Keep exploring the fascinating world of chemistry, and don't hesitate to ask questions. Happy problem-solving!