# AS Level Chemistry Study Notes (2023-2025 Syllabus) ## 1. Atoms, Molecules, and Stoichiometry ### 1.1 Relative Mass - **Atomic Mass ($A_r$):** The weighted average mass of an atom compared to $\text{C}^{12}$, where one atom of $\text{C}^{12}$ is exactly 12 atomic mass units. - **Molecular Mass ($M_r$):** The mass of a molecule, calculated by summing the $A_r$ of all atoms in the molecule. - **Formula Mass:** The mass of one formula unit of a compound (used for ionic compounds). - **Isotopic Mass:** The mass of a specific isotope of an element. - **Unified Atomic Mass Unit ($u$):** Defined as $1.66 \times 10^{-27} \, \text{kg}$. All relative masses are compared to $\text{C}^{12}$, which is the standard reference. --- ### 1.2 The Mole - **Definition of a Mole:** The amount of substance containing the same number of particles (atoms, ions, molecules, or electrons) as there are atoms in exactly 12g of $\text{C}^{12}$. - **Avogadro’s Constant ($N_A$):** $6.02 \times 10^{23}$ particles per mole. 1 mole of water ($\text{H}_2\text{O}$) contains $6.02 \times 10^{23}$ molecules of water. --- ### 1.3 Mass Spectra - **Mass Spectrometry:** Used to determine the relative abundance of isotopes. - **Relative Abundance Formula:** $$ \text{Relative Abundance} = \frac{\text{Peak Height}}{\text{Total Height}} \times 100\% $$ - **Average Atomic Mass Formula:** $$ A_r = \frac{\sum (\text{Mass of Isotope} \times \text{Relative Abundance})}{100} $$ Mass spectra diagrams show peaks corresponding to isotopes. The height of each peak represents the relative abundance. --- ### 1.4 Empirical and Molecular Formulae - **Empirical Formula:** The simplest whole-number ratio of atoms in a compound. - **Molecular Formula:** The actual number of each type of atom in a molecule. #### Key Formula: $$ \text{Molecular Formula} = (\text{Empirical Formula})_n $$ Where: $$ n = \frac{\text{Molecular Mass}}{\text{Mass of Empirical Formula}} $$ A compound has an empirical formula of $\text{CH}_2$ and a molecular mass of 56. Step 1: Calculate the mass of the empirical formula: $12 + (2 \times 1) = 14$. Step 2: Find $n$: $n = \frac{56}{14} = 4$. Step 3: Molecular formula = $\text{C}_4\text{H}_8$. --- ### 1.5 Calculations Involving the Mole Concept #### Key Formulas: 1. **Moles from Mass:** $$ \text{Moles} = \frac{\text{Mass}}{\text{Molar Mass}} $$ 2. **Volume of a Gas at Room Temperature and Pressure (r.t.p.):** $$ \text{Volume (dm}^3\text{)} = \text{Moles} \times 24 $$ 3. **Concentration of a Solution:** $$ \text{Concentration (mol dm}^{-3}\text{)} = \frac{\text{Moles}}{\text{Volume (dm}^3\text{)}} $$ #### Definitions: - **Anhydrous:** A compound with all water molecules removed. - **Hydrated:** A compound containing water molecules in its crystalline structure. - **Water of Crystallization:** The water molecules associated with a hydrated compound. Calculate the moles of $\text{NaCl}$ in 58.5g. Step 1: Molar mass of $\text{NaCl}$ = $23 + 35.5 = 58.5 \, \text{g/mol}$. Step 2: Moles = $\frac{\text{Mass}}{\text{Molar Mass}} = \frac{58.5}{58.5} = 1 \, \text{mol}$. Always ensure units are consistent. For example, convert cm³ to dm³ by dividing by 1000. --- ### Practice Problems 1. Calculate the molecular formula of a compound with an empirical formula of $\text{CH}_3$ and a molecular mass of 30. 2. Determine the volume of $0.5 \, \text{mol}$ of $\text{CO}_2$ gas at r.t.p. 3. A hydrated salt $\text{CuSO}_4 \cdot x\text{H}_2\text{O}$ has a molar mass of 250 g/mol. Find $x$. For gases, always assume standard conditions unless stated otherwise: $1 \, \text{atm}$ pressure and $25^\circ \text{C}$ temperature. --- ### Key Takeaways - The mole is a fundamental concept in chemistry, linking mass, particles, and volume. - Empirical and molecular formulas are essential for understanding the composition of compounds. - Mass spectrometry provides insights into isotopic composition and relative atomic masses. - Always double-check units and conversions in calculations to avoid errors.