Find The Empirical Formula For Tartaric Acid.

Finding the Empirical Formula for Tartaric Acid: A Comprehensive Guide

Determining the empirical formula of a compound is a fundamental concept in chemistry. This process, often performed through combustion analysis or other quantitative techniques, allows us to establish the simplest whole-number ratio of atoms present in a molecule. This article will delve into the process of finding the empirical formula for tartaric acid, a naturally occurring organic acid found in many fruits, particularly grapes. We'll explore the theoretical underpinnings, the practical steps involved, and consider potential sources of error.

Understanding Empirical and Molecular Formulas

Before we embark on the determination of tartaric acid's empirical formula, let's clarify the difference between empirical and molecular formulas.

• Empirical Formula: This represents the simplest whole-number ratio of atoms of each element in a compound. It doesn't necessarily reflect the actual number of atoms present in a single molecule.

• Molecular Formula: This indicates the exact number of atoms of each element present in a single molecule of the compound. It's a multiple of the empirical formula.

For example, the empirical formula for glucose is CH₂O, while its molecular formula is C₆H₁₂O₆. The molecular formula is six times the empirical formula.

The Composition of Tartaric Acid

Tartaric acid is a diprotic acid, meaning it can donate two protons (H⁺ ions). Its chemical name is 2,3-dihydroxybutanedioic acid. This name hints at its composition: it contains carbon (C), hydrogen (H), and oxygen (O) atoms. Our goal is to determine the precise ratio of these atoms in its empirical formula.

Determining the Empirical Formula: A Step-by-Step Approach

The process typically involves the following steps:

1. Combustion Analysis

Combustion analysis is a common technique used to determine the empirical formula of organic compounds. In this method, a precisely weighed sample of the compound is completely burned in an excess of oxygen. The products of this combustion are usually carbon dioxide (CO₂) and water (H₂O). The masses of CO₂ and H₂O produced are carefully measured.

The Crucial Calculations:

• Mass of Carbon (C): The mass of carbon in the CO₂ produced is calculated using the molar mass of C and CO₂. The molar mass of CO₂ is approximately 44 g/mol, and the molar mass of C is 12 g/mol. Therefore, the mass percentage of carbon in CO₂ is (12/44) * 100% ≈ 27.3%.

• Mass of Hydrogen (H): Similarly, the mass of hydrogen in the H₂O produced is determined. The molar mass of H₂O is approximately 18 g/mol, and the molar mass of H is 1 g/mol. Two hydrogen atoms are present in each water molecule, so the mass percentage of hydrogen in H₂O is (2/18) * 100% ≈ 11.1%.

• Mass of Oxygen (O): The mass of oxygen in the original tartaric acid sample can be determined by subtracting the masses of carbon and hydrogen from the total mass of the sample. This assumes that the only elements present are carbon, hydrogen, and oxygen.

2. Calculating Moles

Once the masses of carbon, hydrogen, and oxygen are known, the next step is to convert these masses into moles using their respective molar masses:

• Moles of C = (Mass of C) / (Molar mass of C)
• Moles of H = (Mass of H) / (Molar mass of H)
• Moles of O = (Mass of O) / (Molar mass of O)

3. Determining the Mole Ratio

The next step involves finding the simplest whole-number ratio of moles of each element. This is achieved by dividing the number of moles of each element by the smallest number of moles obtained.

4. Writing the Empirical Formula

Finally, the empirical formula is written using the whole-number ratios obtained in the previous step. The subscripts in the formula represent these whole-number ratios.

Example Calculation

Let's assume a 1.00 g sample of tartaric acid underwent combustion analysis, yielding 1.19 g of CO₂ and 0.40 g of H₂O.

1. Mass of elements:

• Mass of C = (1.19 g CO₂)*(12 g C/44 g CO₂) ≈ 0.325 g C
• Mass of H = (0.40 g H₂O)*(2 g H/18 g H₂O) ≈ 0.044 g H
• Mass of O = 1.00 g - 0.325 g - 0.044 g ≈ 0.631 g O

2. Moles of elements:

• Moles of C = 0.325 g / 12 g/mol ≈ 0.027 mol
• Moles of H = 0.044 g / 1 g/mol ≈ 0.044 mol
• Moles of O = 0.631 g / 16 g/mol ≈ 0.039 mol

3. Mole Ratio:

To find the simplest whole-number ratio, divide each number of moles by the smallest number of moles (0.027 mol):

• C: 0.027 mol / 0.027 mol = 1
• H: 0.044 mol / 0.027 mol ≈ 1.63
• O: 0.039 mol / 0.027 mol ≈ 1.44

These ratios are not whole numbers. To obtain whole numbers, we can multiply by a small integer. In this case, multiplying by 3 gives approximately:

• C: 3
• H: 5
• O: 4

4. Empirical Formula:

Therefore, the empirical formula for tartaric acid based on this example is C₃H₅O₄. This is not the actual empirical formula for tartaric acid, which is C₂H₄O₃. This discrepancy highlights the importance of careful experimental technique and accurate measurements.

Sources of Error

Several factors can introduce errors into the determination of an empirical formula:

• Incomplete Combustion: If the combustion is not complete, the masses of CO₂ and H₂O measured will be lower than expected, leading to inaccurate calculations.

• Impurities in the Sample: The presence of impurities in the tartaric acid sample will affect the results.

• Measurement Errors: Errors in weighing the sample and measuring the masses of CO₂ and H₂O can significantly impact the final result.

• Loss of Products: If some CO₂ or H₂O escapes before being measured, the results will be inaccurate.

Conclusion

Finding the empirical formula of tartaric acid, or any compound, requires meticulous attention to detail and precise measurements. While combustion analysis is a powerful technique, potential sources of error must be considered. The example calculation demonstrates the process, but the result highlights that real-world experiments require careful execution and potential refinement of the method to achieve an accurate result. Remember that the true empirical formula for tartaric acid is C₂H₄O₃. This exercise serves as a learning opportunity to understand the theoretical and practical aspects of determining empirical formulas and analyzing the potential for errors in experimental procedures. The discrepancies between the calculated and actual empirical formulas highlight the sensitivity of this analytical technique and underscores the need for careful execution and potentially repeated trials for reliable results.