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The empirical formula of a compound is \(\mathrm{CH}\). If the molar mass of this compound is about \(78 \mathrm{~g},\) what is its molecular formula?

Short Answer

Expert verified
The molecular formula is \(\mathrm{C}_6\mathrm{H}_6\).

Step by step solution

01

Determine the empirical formula mass

Calculate the molar mass of the empirical formula, CH. The atomic mass of carbon (C) is approximately 12 g/mol and hydrogen (H) is approximately 1 g/mol.\[\text{Empirical formula mass} = (12\, \text{g/mol}) + (1\, \text{g/mol}) = 13\, \text{g/mol}\]
02

Find the ratio of molar mass to empirical formula mass

Divide the given molar mass by the empirical formula mass to find the ratio, which tells you how many times larger the molecular formula is compared to the empirical formula.\[\text{Ratio} = \frac{78\, \text{g/mol}}{13\, \text{g/mol}} = 6\]
03

Determine the molecular formula

Multiply the subscripts in the empirical formula by the ratio obtained to get the molecular formula.Since the empirical formula is \(\text{CH}\) and the ratio is 6, the molecular formula is \(\mathrm{C}_6\mathrm{H}_6\).

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Empirical Formula
The empirical formula provides the simplest whole-number ratio of atoms within a chemical compound. For example, the empirical formula \(\mathrm{CH}\) indicates that carbon and hydrogen exist in a 1:1 ratio in the compound. Unlike the molecular formula, which tells us the exact number of each atom in a molecule, the empirical formula gives a simplified version.
Finding the empirical formula involves determining the relative number of each type of atom by analyzing experimental data, often from mass spectrometry or combustion analysis. It's a fundamental concept that helps chemists understand the composition of compounds, especially when molecular details are complex or unknown.
Molar Mass
The molar mass is the mass of one mole of a substance, and it's measured in grams per mole (g/mol). It's a key concept essential for converting between the amount of substance (in moles) and mass (in grams). Calculating the molar mass requires knowing the atomic masses of all elements in a compound and adding them together based on the compound's formula.
For instance, you would find the molar mass of \(\mathrm{CH}\) by adding the atomic masses of carbon (12 g/mol) and hydrogen (1 g/mol), resulting in a total of 13 g/mol. Understanding molar mass aids in predicting reactions and determining why some reactions might proceed faster than others based on mass interactions.
Chemical Compounds
Chemical compounds consist of two or more different elements joined together by chemical bonds. They can be classified into various types like ionic, covalent, and metallic compounds depending on the nature of the bond. For example, \(\mathrm{C}_6\mathrm{H}_6\), or benzene, is a covalent compound characterized by shared electrons among the atoms.
  • The properties of a compound are distinct from the properties of its constituent elements.
  • Compounds can be broken down through chemical reactions, unlike mixtures which can be separated by physical means.
  • Understanding the composition of compounds is crucial for industries like pharmaceuticals where precision is vital.
Compounds form the backbone of chemistry, acting as reactants, products, and intermediates in countless reactions.
Atomic Mass
Atomic mass is the mass of an individual atom, usually expressed in unified atomic mass units (u or amu). The atomic mass of an element reflects the total number of protons and neutrons in its nucleus. For example, carbon has an atomic mass of approximately 12 amu, while hydrogen is about 1 amu.
Knowing atomic mass is important because it allows chemists to calculate molar masses and form empirical and molecular formulas. This information helps not only in lab settings but also across fields like materials science and biochemistry, where the atomic composition of materials determines their properties and usefulness.
Step by Step Chemistry Solution
A step-by-step solution in chemistry breaks down complex problems into manageable parts. This approach is critical in solving empirical formula and molecular formula problems, among others.
For example, to find the molecular formula from an empirical formula, one first calculates the empirical formula mass. Next, the molar mass of the compound is measured. Then, by dividing the molar mass by the empirical formula mass, you find how many times the empirical formula fits into the molecular formula.
  • This method simplifies problem-solving by reducing potential errors common in holistic approaches.
  • Each step builds understanding, reinforcing the logic behind calculations.
  • This approach is versatile, applicable to other areas such as stoichiometry and reaction balancing.
In essence, step-by-step solutions transform daunting tasks into straightforward, solvable equations through clear, logical progression.

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Most popular questions from this chapter

Aspirin or acetylsalicylic acid is synthesized by combining salicylic acid with acetic anhydride: $$ \mathrm{C}_{7} \mathrm{H}_{6} \mathrm{O}_{3}+\mathrm{C}_{4} \mathrm{H}_{6} \mathrm{O}_{3} \longrightarrow \mathrm{C}_{9} \mathrm{H}_{8} \mathrm{O}_{4}+\mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2} $$ \(\begin{array}{l}\text { salicylic acid acetic anhydride } \\ \text { aspirin } & \text { acetic acid }\end{array}\) (a) How much salicylic acid is required to produce \(0.400 \mathrm{~g}\) of aspirin (about the content in a tablet), assuming acetic anhydride is present in excess? (b) Calculate the amount of salicylic acid needed if only 74.9 percent of salicylic is converted to aspirin. (c) In one experiment, \(9.26 \mathrm{~g}\) of salicylic acid reacts with \(8.54 \mathrm{~g}\) of acetic anhydride. Calculate the theoretical yield of aspirin an the percent yield if only \(10.9 \mathrm{~g}\) of aspirin is produced.

For many years, the extraction of gold - that is, the separation of gold from other materials- involved the use of potassium cyanide: \(4 \mathrm{Au}+8 \mathrm{KCN}+\mathrm{O}_{2}+2 \mathrm{H}_{2} \mathrm{O} \longrightarrow 4 \mathrm{KAu}(\mathrm{CN})_{2}+4 \mathrm{KOH}\) What is the minimum amount of KCN in moles needed to extract \(29.0 \mathrm{~g}\) (about an ounce) of gold?

Monosodium glutamate (MSG), a food-flavor enhancer, has been blamed for "Chinese restaurant syndrome," the symptoms of which are headaches and chest pains. MSG has the following composition by mass: 35.51 percent C \(, 4.77\) percent \(\mathrm{H}, 37.85\) percent \(\mathrm{O}, 8.29\) percent \(\mathrm{N},\) and 13.60 percent Na. What is its molecular formula if its molar mass is about \(169 \mathrm{~g}\) ?

Lactic acid, which consists of \(\mathrm{C}, \mathrm{H},\) and \(\mathrm{O},\) has long been thought to be responsible for muscle soreness following strenuous exercise. Determine the empirical formula of lactic acid given that combustion of a 10.0 -g sample produces \(14.7 \mathrm{~g} \mathrm{CO}_{2}\) and \(6.00 \mathrm{~g} \mathrm{H}_{2} \mathrm{O}\).

Myoglobin stores oxygen for metabolic processes in muscle. Chemical analysis shows that it contains 0.34 percent Fe by mass. What is the molar mass of myoglobin? (There is one Fe atom per molecule.)

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