Determine the simplest formulas of the following compounds: (a) the food enhancer monosodium glutamate (MSG), which has the composition \(35.51 \%\) C, \(4.77 \% \mathrm{H}, 37.85 \% \mathrm{O}, 8.29 \% \mathrm{~N}\), and \(13.60 \% \mathrm{Na} .\) (b) zircon, a diamond-like mineral, which has the composition \(34.91 \%\) \(\mathrm{O}, 15.32 \% \mathrm{Si}\), and \(49.76 \% \mathrm{Zr}\) (c) nicotine, which has the composition \(74.0 \%\) C, \(8.65 \% \mathrm{H}\), and \(17.4 \% \mathrm{~N} .\)

The mole ratio is a foundational concept in the chemistry world, critical to understanding reactions and compounds. It refers to the proportion of moles of one substance compared to the moles of another substance. To find the mole ratio, you need the amount of moles for each element, which you acquire by dividing the mass in grams by the molar mass of that element.
When you're given mass percent composition for a compound, as in our example with monosodium glutamate or other substances, you'll start by assuming a 100g sample. This turns percentages into actual grams, simplifying the use of the molar mass to find moles. Once you've calculated the moles for each element, you divide these by the smallest number of moles you've obtained to find a ratio, eventually leading you to the empirical formula—a simplified version of the compound's molecular formula.
When you're working with an unknown compound's empirical formula, locating the smallest number of moles (the smallest common denominator) is key to finding the mole ratio. This step ensures that you're able to simplify the ratio correctly to reach the most reduced form for the empirical formula.

Mass percent composition is like a recipe for a chemical compound, showing the percentage by mass of each individual element in a compound. It provides a quick look into what's in your chemical 'meal' and how much of each 'ingredient' you have, relative to the total mass of the dish.
This information is indispensable when you're trying to determine the formula of an unknown compound or verifying the purity of a substance. As you saw in our textbook solutions, to determine the empirical formula, we employed the mass percent composition as a starting point. We used the assumption that we have 100 grams of the substance, which conveniently turns the percentage into a direct mass figure in grams. This direct correlation eliminates any need for additional conversion steps and provides a streamlined path to calculating moles, which then leads us to our desired mole ratios and, eventually, the empirical formula.

Stoichiometry is essentially the math behind chemistry, the calculation of quantitative (measurable and numeric) relationships of the substances undergoing chemical reactions. Think of it as an accountant who must balance every debit with a credit; in chemistry, every atom is accounted for.
In the context of finding empirical formulas, stoichiometry includes the process of converting mass to moles, determining the mole ratio, and using this ratio to write the empirical formula. The principles of stoichiometry are grounded on the law of conservation of mass, which states that matter is neither created nor destroyed in chemical reactions, meaning that all atoms present in the reactants must be found in the products.
Hence, stoichiometry not only helps us figure out how much of a reactant is needed to produce a desired amount of product, but it also guides us to the most basic representation of a compound, reflecting the true proportion of elements within.

The term 'molar mass' is frequently tossed around in any chemistry lab. Simply put, it is the weight of one mole (6.022 x 10²³ particles) of a given substance. It is measured in grams per mole (g/mol) and can be found on the periodic table as the atomic weight of an element.
In the exercises we've discussed, the molar mass serves as a conversion factor between grams of a substance and the number of moles of that substance. It's vital in chemical calculations, as it bridges the gap between the microscopic world of atoms and molecules and the macroscopic world where we measure things like mass and volume. Understanding how to use molar mass in calculations allows you to take a percent composition by mass and turn it into a mole ratio, the first and most crucial step in determining the compound's empirical formula.