Unit 3: Comprehensive Summary

Practical chemistry requires a strict adherence to safety protocols to prevent accidents. This includes wearing Personal Protective Equipment (PPE) like goggles and lab coats, proper handling of chemicals, and knowing the location of emergency equipment.

All measurements are based on base quantities (e.g., mass, time, amount of substance) and derived quantities (e.g., volume, density). Accurate measurements depend on selecting the correct instrument:

• Volumetric Pipettes and Flasks: Used for high-precision, fixed volume measurements, essential for preparing standard solutions.
• Burettes: Used for delivering variable, precise volumes, primarily in titrations.
• Measuring Cylinders: Used for approximate volume measurements.
• Analytical Balances: Used for highly precise mass measurements.

Understanding the difference between accuracy (closeness to the true value) and precision (consistency of results) is key to evaluating experimental data.

The mole is the SI unit for the amount of a substance. One mole contains exactly Avogadro's number ($N_A$) of particles, which is $6.022 \times 10^{23}$ particles. This constant allows us to connect the microscopic world of atoms and molecules to the macroscopic world of grams and liters.

The molar mass (M) of a substance is the mass of one mole, expressed in grams per mole (g/mol). It is numerically equal to the relative molecular mass ($M_r$) for covalent compounds or the relative formula mass for ionic compounds. The fundamental relationship connecting mass (m), moles (n), and molar mass (M) is: $$ n = \frac{m}{M} $$

For solutions, we use molarity (C) to express concentration in moles of solute per liter (or dm³) of solution. The relationship is: $n = C \times V$. When diluting a stock solution, the moles of solute remain constant, leading to the dilution formula: $M_1V_1 = M_2V_2$.

For gases, the relationship between pressure (P), volume (V), moles (n), and temperature (T) is described by the Ideal Gas Equation: $$ PV = nRT $$ Where R is the molar gas constant ($8.31 \, J \, mol^{-1} K^{-1}$). For quick calculations under standard conditions, we can use the molar volume of a gas:

• At STP (Standard Temperature and Pressure: 273 K, 101 kPa), 1 mole of any gas occupies 22.4 dm³.
• At RTP (Room Temperature and Pressure: 298 K, 101 kPa), 1 mole of any gas occupies 24.4 dm³.

The empirical formula represents the simplest whole-number ratio of atoms in a compound (e.g., $CH_2O$ for glucose). The molecular formula shows the actual number of atoms of each element in a molecule (e.g., $C_6H_{12}O_6$ for glucose).

To find the empirical formula from percentage composition, we convert mass to moles, find the simplest mole ratio, and convert it to whole numbers. To find the molecular formula, we first determine the empirical formula mass and then find the whole-number multiple 'n' by which it relates to the given molecular mass: $$ n = \frac{\text{Molecular Mass}}{\text{Empirical Formula Mass}} $$

A chemical equation is a symbolic representation of a chemical reaction. It must be balanced to satisfy the Law of Conservation of Mass, ensuring the same number of each type of atom on both the reactant and product sides. The numbers used to balance equations are called stoichiometric coefficients.

Stoichiometry is the process of using a balanced chemical equation to calculate the quantitative relationships between reactants and products. The core of all stoichiometric calculations is the mole ratio, which acts as a conversion factor to relate the moles of any two substances in the reaction. The general process is:

1. Convert the given quantity (mass, volume, etc.) into moles.
2. Use the mole ratio from the balanced equation to find the moles of the desired substance.
3. Convert the moles of the desired substance back into the required units (mass, volume, etc.).