Determine Molecular Formula Compound Contains Carbon And Hydrogen

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In the realm of chemistry, determining the molecular formula of a compound is a fundamental yet crucial task. The molecular formula provides a precise representation of the number and types of atoms present in a molecule, offering invaluable insights into its structure, properties, and reactivity. This article delves into a step-by-step approach to deduce the molecular formula of a hydrocarbon compound, given its elemental composition and molar mass. We will embark on a journey through calculations and reasoning, ultimately unveiling the true identity of the compound.

At the heart of molecular formula determination lies the concept of the empirical formula. The empirical formula represents the simplest whole-number ratio of atoms in a compound. It serves as a crucial stepping stone towards deciphering the molecular formula. To calculate the empirical formula, we must first convert the given mass percentages of each element into moles. This conversion involves dividing the mass of each element by its respective atomic mass, which can be readily obtained from the periodic table.

In this specific scenario, we are presented with a sample containing 60.0 grams of carbon (C) and 5.05 grams of hydrogen (H). To convert these masses into moles, we divide them by the atomic masses of carbon (12.01 g/mol) and hydrogen (1.01 g/mol), respectively. This yields approximately 5 moles of carbon and 5 moles of hydrogen. The ratio of moles of carbon to hydrogen is thus 1:1. This signifies that the simplest whole-number ratio of carbon to hydrogen atoms in the compound is 1:1. Therefore, the empirical formula of the compound is CH.

With the empirical formula in hand, the next step is to determine the molecular formula. The molecular formula represents the actual number of atoms of each element present in a molecule of the compound. It is a whole-number multiple of the empirical formula. To establish this multiple, we compare the molar mass of the compound with the molar mass of the empirical formula. The ratio of these molar masses provides the factor by which the empirical formula must be multiplied to obtain the molecular formula.

In our case, the molar mass of the compound is given as 78.12 g/mol. The molar mass of the empirical formula (CH) can be calculated by summing the atomic masses of carbon and hydrogen, which amounts to approximately 13.02 g/mol. Dividing the molar mass of the compound (78.12 g/mol) by the molar mass of the empirical formula (13.02 g/mol) yields a ratio of approximately 6. This indicates that the molecular formula is six times the empirical formula.

Multiplying the subscripts in the empirical formula (CH) by the factor of 6, we arrive at the molecular formula of the compound: C6H6. This formula reveals that each molecule of the compound contains 6 carbon atoms and 6 hydrogen atoms. Among the options provided, C6H6 is the correct molecular formula.

This methodical approach of determining the empirical formula and subsequently using it to derive the molecular formula is a cornerstone of chemical analysis. It enables chemists to unravel the composition of unknown compounds and gain deeper insights into their molecular structures and properties. By understanding the fundamental principles of stoichiometry and molar mass calculations, we can confidently navigate the intricate world of molecular formulas and unlock the secrets hidden within chemical compounds.

To solidify our understanding of molecular formula determination, let's delve deeper into the underlying principles and explore some nuances that can arise in more complex scenarios. We will also address some frequently asked questions to provide a comprehensive perspective on this crucial aspect of chemistry.

Understanding the Significance of Molar Mass

Molar mass, the mass of one mole of a substance, plays a pivotal role in determining the molecular formula. It acts as a bridge connecting the macroscopic world of grams and the microscopic world of atoms and molecules. The molar mass is numerically equivalent to the sum of the atomic masses of all the atoms in a molecule, expressed in grams per mole (g/mol). This fundamental concept allows us to relate the mass of a substance to the number of molecules it contains.

The molar mass of a compound can be experimentally determined using various techniques, such as mass spectrometry or colligative properties measurements. Once the molar mass is known, it serves as a crucial piece of information in deciphering the molecular formula.

Dealing with Non-Integer Ratios

In some instances, the mole ratios calculated from the elemental composition may not be whole numbers. This can occur due to experimental errors or the presence of isotopes. To obtain the empirical formula, we need to convert these non-integer ratios into whole numbers. This is typically achieved by dividing all the mole ratios by the smallest mole ratio. If the resulting ratios are still not whole numbers, we may need to multiply all the ratios by a common factor to obtain whole-number subscripts.

For example, if the mole ratios of carbon, hydrogen, and oxygen in a compound are 1.5, 2.25, and 1, respectively, we would first divide all ratios by the smallest ratio (1), resulting in 1.5, 2.25, and 1. To convert these to whole numbers, we can multiply all ratios by 4, yielding the whole-number ratios of 6, 9, and 4. The empirical formula would then be C6H9O4.

The Importance of Experimental Accuracy

The accuracy of the determined molecular formula hinges on the precision of the experimental data used in the calculations. Errors in mass measurements or molar mass determination can lead to incorrect mole ratios and, consequently, an erroneous molecular formula. Therefore, it is imperative to employ accurate analytical techniques and exercise meticulous care in data collection.

Moreover, the purity of the sample being analyzed is crucial. The presence of impurities can distort the elemental composition and lead to inaccurate results. Sample purification techniques, such as recrystallization or distillation, may be necessary to ensure the integrity of the data.

To reinforce your understanding of molecular formula determination, let's work through some practice problems.

Problem 1: A compound contains 40.0% carbon, 6.7% hydrogen, and 53.3% oxygen by mass. Its molar mass is 180.16 g/mol. Determine the compound's molecular formula.

Solution:

  1. Convert percentages to grams: Assume 100 g of the compound. This gives 40.0 g C, 6.7 g H, and 53.3 g O.
  2. Convert grams to moles:
    • Moles of C = 40.0 g / 12.01 g/mol = 3.33 mol
    • Moles of H = 6.7 g / 1.01 g/mol = 6.63 mol
    • Moles of O = 53.3 g / 16.00 g/mol = 3.33 mol
  3. Determine the mole ratio: Divide each mole value by the smallest (3.33):
    • C: 3.33 / 3.33 = 1
    • H: 6.63 / 3.33 ≈ 2
    • O: 3.33 / 3.33 = 1
  4. The empirical formula is CH2O.
  5. Calculate the empirical formula mass: 12.01 + 2(1.01) + 16.00 = 30.03 g/mol
  6. Determine the multiplier: 180.16 g/mol / 30.03 g/mol ≈ 6
  7. Multiply the subscripts in the empirical formula by 6 to get the molecular formula: C6H12O6

Problem 2: A compound has the following elemental composition: 24.24% C, 4.07% H, and 71.65% Cl. The molar mass of the compound is 98.96 g/mol. What is the molecular formula?

Solution:

  1. Convert percentages to grams (assuming 100 g sample): 24.24 g C, 4.07 g H, 71.65 g Cl
  2. Convert grams to moles:
    • Moles C = 24.24 g / 12.01 g/mol ≈ 2.02 mol
    • Moles H = 4.07 g / 1.01 g/mol ≈ 4.03 mol
    • Moles Cl = 71.65 g / 35.45 g/mol ≈ 2.02 mol
  3. Divide each mole quantity by the smallest (2.02) to find the mole ratio:
    • C: 2.02 / 2.02 = 1
    • H: 4.03 / 2.02 ≈ 2
    • Cl: 2.02 / 2.02 = 1
  4. The empirical formula is CH2Cl.
  5. Calculate the empirical formula mass: 12.01 + 2(1.01) + 35.45 ≈ 49.48 g/mol
  6. Divide the molar mass by the empirical formula mass: 98.96 g/mol / 49.48 g/mol ≈ 2
  7. Multiply the subscripts in the empirical formula by 2 to obtain the molecular formula: C2H4Cl2

These practice problems illustrate the step-by-step process of determining molecular formulas. By mastering these techniques, you can confidently tackle a wide range of chemical problems involving compound composition and structure.

The determination of molecular formulas is a cornerstone of chemical analysis, providing invaluable insights into the composition and structure of compounds. By systematically converting elemental composition data into mole ratios, establishing the empirical formula, and leveraging molar mass information, we can successfully unveil the molecular formula. The molecular formula, in turn, serves as a foundation for understanding a compound's properties, reactivity, and role in chemical reactions. Continuous practice and a deep understanding of the underlying principles will empower you to excel in this essential area of chemistry.

To further enhance your understanding, let's address some frequently asked questions related to molecular formula determination:

Q1: What is the difference between an empirical formula and a molecular formula?

  • The empirical formula represents the simplest whole-number ratio of atoms in a compound, while the molecular formula represents the actual number of atoms of each element in a molecule.

Q2: Why is it necessary to convert mass percentages to moles when determining empirical formulas?

  • Moles provide a direct comparison of the number of atoms of each element in a compound, as one mole of any element contains the same number of atoms (Avogadro's number). Mass percentages, on the other hand, do not directly reflect the atomic ratios due to differences in atomic masses.

Q3: What happens if the mole ratios are not whole numbers?

  • Divide all mole ratios by the smallest mole ratio. If the resulting ratios are still not whole numbers, multiply all ratios by a common factor to obtain whole-number subscripts for the empirical formula.

Q4: How does the molar mass help in determining the molecular formula?

  • The molar mass of the compound is divided by the molar mass of the empirical formula to find the factor by which the empirical formula must be multiplied to obtain the molecular formula.

Q5: Can two different compounds have the same empirical formula?

  • Yes, different compounds can have the same empirical formula if they have the same simplest whole-number ratio of atoms. However, they will have different molecular formulas and molar masses.

Q6: What are some common sources of error in molecular formula determination?

  • Errors in mass measurements, inaccurate molar mass determination, and the presence of impurities in the sample can lead to errors in the calculated molecular formula. Accurate analytical techniques and sample purification are crucial for reliable results.

Q7: Is it always possible to determine the molecular formula from the empirical formula?

  • No, it is not always possible. To determine the molecular formula, you need the molar mass of the compound in addition to the empirical formula. The molar mass provides the information needed to determine the multiple by which the empirical formula must be scaled to obtain the molecular formula.

By addressing these FAQs, we aim to provide a comprehensive understanding of molecular formula determination and its practical applications in chemistry.

A. CH B. $C _2 H _2$ C. $C _6 H _6$ D. $C _6 H$

Correct answer is C. $C _6 H _6$

Molecular Formula, Empirical Formula, Stoichiometry, Molar Mass, Chemical Composition, Hydrocarbon, Mole Ratio, Elemental Analysis, Chemical Calculation, Chemistry Problem