Basic Principles of Chemistry and Stoichiometry Study Notes

Chemistry - The Centre of Life

• The Significance of Chemistry in Basic Human Needs: In the Tamil classical language, the phrase 'Unna unavu, udukka udai, irukka idam' highlights the three fundamental needs: food, clothing, and shelter. Chemistry plays a vital role in providing these through:

  • Agriculture: Production of fertilizers and insecticides to enhance output.

  • Construction: Developing modern cements, concrete mixtures, and high-quality steel for weather-resistant buildings.

  • Textiles: Creating better quality fabrics.

• The Ubiquity of Chemistry: Chemistry is present everywhere. The human body is composed of chemicals and functions via continuous biochemical reactions. It impacts culture, life, and the environment.

• Modern Challenges and Evolution: Chemistry expands to meet modern challenges. Chemical industries produce polymers, dyes, alloys, and life-saving drugs.

• Case Study: HIV/AIDS: In the early 1980s, HIV/AIDS patients rarely survived beyond a few years. Chemistry has since provided effective medicines, allowing infected individuals to live longer, better lives.

• Environmental Chemistry: Chemical principles led to the replacement of ozone-depleting CFCs (Chlorofluorocarbons) in refrigerators with eco-friendly alternatives and the development of "green processes."

Classification of Matter

• Definition of Matter: Matter is defined as anything that has mass and occupies space. All matter is composed of atoms.

• Methods of Classification:

  • Physical State: Matter is classified as solids, liquids, or gases. These states can be interconverted by adjusting temperature and pressure.

  • Chemical Composition: Matter is divided into mixtures and pure substances.

• Mixtures: Consist of multiple chemical entities without chemical interaction.

  • Homogeneous Mixtures: Uniform physical appearance (e.g., Green tea).

  • Heterogeneous Mixtures: Non-uniform physical appearance (e.g., Oil and water).

• Pure Substances: Composed of simple atoms or molecules.

  • Elements: Consist of only one type of atom. Atoms are the smallest electrically neutral particles made of electrons, protons, and neutrons.

    • Monatomic Units: Gold ($Au$), Copper ($Cu$).

    • Polyatomic Units (Molecules): Hydrogen ($H_2$), Phosphorus ($P_4$), Sulphur ($S_8$).

  • Compounds: Made of molecules containing two or more different elements. Properties of a compound differ from its constituent elements. For example, Sodium ($Na$) is a shiny metal and Chlorine ($Cl_2$) is an irritating gas, but their compound Sodium Chloride ($NaCl$) is a crystalline solid vital for life.

Atomic and Molecular Masses

• The Scale of Atoms: Individual atoms are too small to measure directly. They have diameters of approximately $10^{-10}\,\text{m}$ and weigh approximately $10^{-27}\,\text{kg}$.

• Standard Scale: IUPAC uses Carbon-12 ($^{12}C$) in its ground state as the standard, fixing its mass at $12\,u$ (unified atomic mass units).

• Unified Atomic Mass Unit ($u$): Defined as one-twelfth of the mass of a Carbon-12 atom.

  • $1\,\text{amu (or) } 1\,u = 1.6605 \times 10^{-27}\,\text{kg}$

• Relative Atomic Mass ($A_r$): The ratio of the average atomic mass to the unified atomic mass unit.

  • Example for Hydrogen ($A_r$)H: $\frac{1.6736 \times 10^{-27}\,\text{kg}}{1.6605 \times 10^{-27}\,\text{kg}} = 1.0078 \approx 1.008\,u$

• Average Atomic Mass: Calculated based on the natural abundance of isotopes.

  • Example: Chlorine naturally occurs as $^{35}Cl$ and $^{37}Cl$ in a ratio of $77:23$.

  • Average relative atomic mass = $\frac{(35 \times 77) + (37 \times 23)}{100} = 35.46\,u$

• Relative Molecular Mass: Calculated by adding the relative atomic masses of all constituent atoms in a molecule.

  • Hydrogen molecule ($H_2$): $2 \times 1.008\,u = 2.016\,u$

  • Glucose ($C_6H_{12}O_6$): $(6 \times 12) + (12 \times 1.008) + (6 \times 16) = 180.096\,u$

Mole Concept and Molar Mass

• The Mole: The SI unit representing a specific amount of substance. Analogous to a "dozen" (12) or a "quire" (25).

• Definition: One mole is the amount of substance containing as many elementary particles (atoms, molecules, ions, etc.) as there are atoms in $12\,\text{g}$ of Carbon-12.

• Avogadro Number ($N_A$): The number of entities in one mole is $6.022 \times 10^{23}$. This is named after Amedeo Avogadro.

• Calculations involving One Mole:

  • $12\,\text{g}$ of Carbon contains $6.022 \times 10^{23}$ atoms.

  • $158.03\,\text{g}$ of Potassium Permanganate ($KMnO_4$) contains $6.022 \times 10^{23}$ molecules.

  • $294.18\,\text{g}$ of Potassium Dichromate ($K_2Cr_2O_7$) contains $6.022 \times 10^{23}$ molecules.

• Molar Mass: The mass of one mole of a substance, expressed in $\text{g mol}^{-1}$. It is numerically equal to the relative atomic/molecular mass.

• Molar Volume: The volume occupied by one mole of a gas at specific conditions:

  • $273\,\text{K}$ and $1\,\text{bar}$ (STP): $22.71\,\text{L}$

  • $273\,\text{K}$ and $1\,\text{atm}$ (SATP): $22.4\,\text{L}$

  • $298\,\text{K}$ and $1\,\text{atm}$ (Room Temp): $24.5\,\text{L}$

Gram Equivalent Concept

• Definition: Gram equivalent mass is the mass of an entity that combines with or displaces $1.008\,\text{g}$ Hydrogen, $8\,\text{g}$ Oxygen, or $35.5\,\text{g}$ Chlorine.

• Formula: $\text{Gram equivalent mass} = \frac{\text{Molar mass (g mol}^{-1})}{\text{Equivalence factor (eq mol}^{-1})}$

• Equivalence Factors ($n$):

  • Acids: Basicity (number of ionisable $H^{+}$ ions). For $H_2SO_4$, $n=2$. Equivalent mass = $\frac{98}{2} = 49\,\text{g eq}^{-1}$.

  • Bases: Acidity (number of ionisable $OH^{-}$ ions). For $KOH$, $n=1$. Equivalent mass = $\frac{56}{1} = 56\,\text{g eq}^{-1}$.

  • Oxidising/Reducing Agents: Number of moles of electrons gained or lost. For $KMnO_4$ in acid medium, $n=5$ ($MnO_4^{-} + 8H^{+} + 5e^{-} \rightarrow Mn^{2+} + 4H_2O$). Equivalent mass = $\frac{158}{5} = 31.6\,\text{g eq}^{-1}$.

Empirical and Molecular Formulas

• Empirical Formula: The simplest whole-number ratio of atoms in a compound.

• Molecular Formula: The actual number of atoms of each element in a molecule. It is a whole-number multiple of the empirical formula ($\text{Molecular Formula} = n \times \text{Empirical Formula}$).

• Determining Empirical Formula:

  1. Assume $100\,\text{g}$ total mass (percentages become grams).

  2. Divide mass by atomic mass to get relative number of moles.

  3. Divide each by the smallest molar value to get the simplest ratio.

  4. Convert to whole numbers if necessary.

• Whole Number Calculation: $n = \frac{\text{Molar mass}}{\text{Calculated empirical formula mass}}$.

  • Example: Vinegar (Acetic acid) and Sour milk (Lactic acid) both have empirical formula $CH_2O$ (mass = 30). Acetic acid (molar mass 60, $n=2$) is $C_2H_4O_2$. Lactic acid (molar mass 90, $n=3$) is $C_3H_6O_3$.

Stoichiometry and Limiting Reagents

• Stoichiometry: Quantity relationship based on balanced equations (Greek: stoicheion = element, metron = measure).

• Limiting Reagent: The reactant that is completely consumed in a reaction, limiting the amount of product formed.

• Excess Reagent: The reactant that remains after the reaction stops.

• Calculations: Stoichiometric coefficients represent the mole-to-mole ratio. These can be used to convert between mass, moles, and volume ($1\,\text{mole gas} = 22.4\,\text{L at STP}$).

Redox Reactions and Oxidation Numbers

• Classical Definitions:

  • Oxidation: Addition of oxygen or removal of hydrogen.

  • Reduction: Removal of oxygen or addition of hydrogen.

• Electronic Concept:

  • Oxidation: Loss of electrons ($Fe^{2+} \rightarrow Fe^{3+} + e^{-}$).

  • Reduction: Gain of electrons ($Cu^{2+} + 2e^{-} \rightarrow Cu$).

• Oxidation Number Rules:

  1. Free elements = 0.

  2. Monatomic ions = charge of ion.

  3. Sum of oxidation states in a neutral molecule = 0; in ions, it equals the net charge.

  4. Hydrogen = $+1$ (except in metal hydrides, where it is $-1$).

  5. Fluorine = $-1$ always.

  6. Oxygen = $-2$ (Exceptions: Peroxides $-1$, Superoxides $-\frac{1}{2}$, with Fluorine ($OF_2$) $+2$).

  7. Alkali metals = $+1$, Alkaline earth metals = $+2$.

• Types of Redox Reactions*:

  1. Combination: $C + O_2 \rightarrow CO_2$.

  2. Decomposition: $2KClO_3 \rightarrow 2KCl + 3O_2$.

  3. Displacement:

     • Metal: $CuSO_4 + Zn \rightarrow Cu + ZnSO_4$.

     • Non-Metal: $Zn + 2HCl \rightarrow ZnCl_2 + H_2$.

  4. Disproportionation: Same element is both oxidised and reduced ($2H_2O_2 \rightarrow 2H_2O + O_2$).

  5. Competitive Electron Transfer: Metals vary in electron-releasing tendency. Order: Zn > Cu > Ag.

Balancing Redox Reactions

• Oxidation Number Method: Steps involve identifying atoms changing oxidation state, balancing electrons through cross-multiplication, and then balancing other atoms ($H$ and $O$ last).

• Ion-Electron (Half-Reaction) Method: Reaction is split into oxidation and reduction halves. Atoms and charges are balanced separately for each half before recombining.

Questions & Discussion

• Evaluate Yourself Exercise 1: Classify Sugar (compound), Sea water (mixture), Distilled water (compound), Carbon dioxide (compound), Copper wire (element), Table salt (compound), Silver plate (element), Naphthalene balls (compound).

• Antacids and Heartburn: Gastric acid ($HCl$) typically has a concentration of $0.082\,\text{M}$. Heartburn occurs when it exceeds $0.1\,\text{M}$. Antacids like $Al(OH)_3$ and $Mg(OH)_2$ neutralise the acid.

  • $3HCl + Al(OH)_3 \rightarrow AlCl_3 + 3H_2O$

  • $2HCl + Mg(OH)_2 \rightarrow MgCl_2 + 2H_2O$

  • Calculation: One tablet with $250\,\text{mg}$ each of $Al(OH)_3$ and $Mg(OH)_2$ neutralises $0.0182\,\text{mole of HCl}$.

• Haemoglobin and Oxygen Transport: Iron ($Fe^{2+}$) in haemoglobin binds oxygen in an octahedral geometry. It is protected from oxidation to $Fe^{3+}$ (Methemoglobin) by the globin protein's hydrophobic environment. About $3\%$ of haemoglobin is oxidized daily but reduced back by the enzyme methemoglobin reductase.

• Cyanide Poisoning: Cyanide binds irreversibly to haemoglobin, blocking oxygen transport and leading to rapid death.