Which Of The Following Compounds Have The Same Empirical Formula

Deciphering Empirical Formulas: Exploring Chemical Compounds with Identical Simplifications

Understanding empirical formulas is fundamental to chemistry. An empirical formula represents the simplest whole-number ratio of atoms in a compound. It doesn't necessarily reflect the actual number of atoms present in a molecule (that's the molecular formula), but rather the ratio between them. This means several different compounds can share the same empirical formula. This article will delve deep into this concept, exploring various examples and explaining how to determine if compounds possess the same empirical formula.

What is an Empirical Formula?

Before we dive into comparing compounds, let's solidify our understanding of empirical formulas. It's the most reduced form of a chemical formula. For instance, consider glucose (C₆H₁₂O₆). Its empirical formula is CH₂O, representing the simplest ratio of carbon, hydrogen, and oxygen atoms (1:2:1). This simplification is crucial in analytical chemistry, where determining the exact molecular formula might be challenging, but the relative proportions of elements can be readily obtained through techniques like elemental analysis.

Key Differences between Empirical and Molecular Formulas:

• Empirical Formula: Shows the simplest whole-number ratio of atoms in a compound.
• Molecular Formula: Shows the actual number of atoms of each element in a molecule of the compound.

The molecular formula is always a whole-number multiple of the empirical formula. For glucose, the molecular formula (C₆H₁₂O₆) is six times the empirical formula (CH₂O).

Determining Empirical Formulas: A Step-by-Step Guide

Let's outline a practical method for determining the empirical formula of a compound, a crucial first step in comparing compounds for identical empirical formulas. Assume we have the percentage composition of a compound:

1. Convert percentages to grams: Assume you have a 100-gram sample. Percentages directly translate to grams in this scenario. For example, 40% carbon becomes 40 grams of carbon.

2. Convert grams to moles: Use the molar mass of each element (found on the periodic table) to convert the grams of each element into moles. For instance, for carbon, you'd divide the grams of carbon by its molar mass (approximately 12.01 g/mol).

3. Find the mole ratio: Divide the number of moles of each element by the smallest number of moles obtained in step 2. This provides the simplest whole-number ratio of atoms.

4. Express the empirical formula: Use the whole-number ratios obtained in step 3 as subscripts for each element in the empirical formula.

Example:

Let's determine the empirical formula of a compound containing 40% carbon, 6.7% hydrogen, and 53.3% oxygen.

1. Grams: 40 g C, 6.7 g H, 53.3 g O
2. Moles:
   • Carbon: 40 g / 12.01 g/mol ≈ 3.33 mol
   • Hydrogen: 6.7 g / 1.01 g/mol ≈ 6.63 mol
   • Oxygen: 53.3 g / 16.00 g/mol ≈ 3.33 mol
3. Mole Ratio:
   • Carbon: 3.33 mol / 3.33 mol = 1
   • Hydrogen: 6.63 mol / 3.33 mol ≈ 2
   • Oxygen: 3.33 mol / 3.33 mol = 1
4. Empirical Formula: CH₂O

Comparing Compounds for Identical Empirical Formulas: Case Studies

Now, let's examine several compounds and determine which share the same empirical formula. We'll use various examples to highlight different scenarios and complexities.

Case Study 1: Simple Hydrocarbons

• Ethene (C₂H₄): Its empirical formula is CH₂.
• Propene (C₃H₆): Its empirical formula is CH₂.
• Butene (C₄H₈): Its empirical formula is CH₂.

In this case, ethene, propene, and butene all have the same empirical formula, CH₂, even though their molecular formulas differ. This demonstrates that multiple compounds can share the same simplest whole-number ratio of atoms.

Case Study 2: Compounds with Oxygen

• Formaldehyde (CH₂O): Empirical formula is CH₂O.
• Acetic acid (C₂H₄O₂): Empirical formula is CH₂O.
• Glucose (C₆H₁₂O₆): Empirical formula is CH₂O.

Here, formaldehyde, acetic acid, and glucose all share the empirical formula CH₂O, showcasing that the ratio of atoms can be the same despite significant variations in the molecular structure and size.

Case Study 3: Incorporating Different Elements

• Hydrogen Peroxide (H₂O₂): Empirical formula is HO.
• Water (H₂O): Empirical formula is H₂O.

While both contain hydrogen and oxygen, their empirical formulas are distinct. This highlights the importance of precise calculation in determining empirical formulas.

Case Study 4: Dealing with Fractional Ratios

Sometimes, the initial mole ratio isn't a whole number. In this situation, you must multiply all the ratios by a factor to obtain whole numbers. For example, if you get a ratio of 1:1.5, you would multiply both by 2 to get 2:3.

Advanced Considerations and Applications

The concept of empirical formulas extends far beyond simple comparisons. It plays a crucial role in several areas of chemistry:

• Elemental Analysis: Determining the empirical formula is a common step in identifying unknown compounds using techniques such as combustion analysis.

• Polymer Chemistry: Understanding the empirical formula of a repeating unit is essential for characterizing polymers.

• Stoichiometry: Empirical formulas are used in stoichiometric calculations to determine the amounts of reactants and products involved in chemical reactions.

Conclusion: The Significance of Empirical Formulas

Determining and comparing empirical formulas is a fundamental skill in chemistry. While the molecular formula provides complete information about the composition of a molecule, the empirical formula offers a simplified representation, revealing the simplest ratio of atoms present. This simplified view allows for comparisons between compounds, revealing unexpected similarities and offering insights into their chemical relationships. Mastering the process of calculating and interpreting empirical formulas unlocks a deeper understanding of the world of chemical compounds and their interrelationships. Understanding the distinction between empirical and molecular formulas is crucial for accurate chemical analysis and interpretation. Remember that multiple compounds can share the same empirical formula, highlighting the importance of considering both empirical and molecular formulas to fully characterize a substance.