Molecules and Stoichiometry for the DAT

Learn key DAT concepts related to moles, molecules, and stoichiometry, plus practice questions and answers

learn everything about molecules and stoichiometry for the dat

Table of Contents

Part 1: Introduction to molecules and stoichiometry

Part 2: Moles

a) Definition of a mole

b) Molar mass and density

Part 3: Formation of molecules

a) Molecular and empirical formulas

b) Types of chemical reactions

Part 4: Stoichiometry

a) Balancing equations

b) Determining theoretical yield

Part 5: High-yield terms

Part 6: Questions and answers

----

Part 1: Introduction to molecules and stoichiometry

Molecules are a fundamental concept in chemistry, and everything we interact with is made of them. A molecule is a group of two or more atoms that are bonded together. For example, salt molecules are made up of sodium and chloride, while water molecules contain two hydrogen atoms and one oxygen atom. Molecules can be hard to visualize and understand, however, because we can’t see them with the naked eye. Even some of the tiniest quantities we can see, such as a grain of salt or drop of water, contain more than one quintillion (10¹⁸) molecules.

Stoichiometry is a way to understand and quantify molecules and atoms. As you study and practice stoichiometry, try to keep this purpose in mind.

Throughout this guide, pay attention to any bold terms, and use the practice questions at the end to apply what you’re learning.

----

Part 2: Moles

a) Definition of a mole

How do we measure the quantities of reactants and products? The quantity of a substance can be expressed in terms of mass, the number of particles, or other relevant units.

Counting the number of particles (atoms or molecules) within a quantity of reactants or products can be challenging due to their extremely small size. (For a comprehensive overview of atomic and molecular structure, please refer to our guide on the topic.) To define these quantities, we use a unit called a "mole".

The mole (abbreviated as "mol") is a versatile unit that allows us to relate the number of atoms or molecules to their mass in grams. Specifically, Avogadro’s number (N_A) quantifies the number of particles in one mole, with N_A = 6.022 x 10²³. When you have the number of particles, Avogadro’s number (N_A) serves as a conversion factor, enabling you to convert between the number of atoms or molecules and the number of moles, or vice versa.

FIGURE 1: DIFFERENT WAYS OF EXPRESSING THE SAME QUANTITY OF A SUBSTANCE.

A mole can also be defined when dealing with gaseous substances. Specifically, one mole of a gaseous compound occupies a volume of 22.4 liters under standard temperature and pressure (STP) conditions. Standard temperature and pressure can be defined as a temperature of 273.15 Kelvin (K) and an absolute pressure of 100 kilopascals (kPa), which is equivalent to 1 bar. Note that this temperature is not equal to room temperature (approx. 293 K).

b) Molar mass and density

The atomic or molecular mass of a substance is a crucial factor for converting the number of moles present into the corresponding mass in grams. This property is also commonly known as molar mass or molecular weight.

Determining the molar mass of an atomic element is a relatively straightforward process. In the case of monatomic elements, the molar mass is the same as the atomic mass, which can be found directly below the element’s symbol on the periodic table. For example, sodium (Na) has an atomic mass of 22.99 atomic mass units (AMU). Since 1 AMU is equivalent to 1 gram per mole (g/mol), this number enables us to convert the mass of a known quantity of sodium into the number of moles in that sample.

FIGURE 2: CONVERTING 15.00 G OF SODIUM (NA) TO THE CORRESPONDING QUANTITY OF MOLES.

Certain elements exist as diatomic molecules, meaning they are composed of two atoms bonded together rather than existing as single atoms. There are seven diatomic elements, and you should remember each of them. These elements are hydrogen (H₂), nitrogen (N₂), oxygen (O₂), and the halogens (F₂, Cl₂, I₂, Br₂).

To calculate the molar mass of these diatomic elements, we need to multiply the atomic mass of the element by the number of atoms in the bonded molecule. Let's consider oxygen as an example. The atomic mass of a single oxygen atom is approximately 16 atomic mass units (AMU). However, in its natural state, oxygen exists as a diatomic molecule (O₂). To determine its molar mass, we must multiply this atomic mass by 2 because there are two oxygen atoms present in each O₂ molecule.

FIGURE 3: CONVERTING 15.00 G OF DIATOMIC OXYGEN (O2) TO THE CORRESPONDING QUANTITY OF MOLES.

When calculating the molar mass of molecular compounds, it's crucial to consider the identity and number of the atoms within a single molecule of the compound. Let's calculate the molar mass of Na₂SO₄ (sodium sulfate) as an example.