Some Basic Concepts of Chemistry – Comprehensive Notes

This foundational chapter is the gateway to understanding the entire realm of chemistry. It lays down the essential definitions, laws, and quantitative relationships crucial for mastering inorganic, organic, and physical chemistry, making it indispensable for NEET and JEE Main aspirants.

1. Importance and Scope of Chemistry

• Chemistry Defined: The branch of science that deals with the composition, structure, properties, and changes of matter. It also explores the energy changes accompanying these transformations.
• Central Science: Chemistry acts as a bridge between physics (at a fundamental level) and biology, geology, environmental science, and materials science, impacting almost every aspect of our lives.
• Applications: It underpins advancements in medicine (drug development), food production and preservation, development of new materials (polymers, composites), energy generation (fuels, batteries), environmental sustainability (pollution control, green chemistry), and industrial processes.

2. Nature of Matter

• Matter: Anything that possesses mass and occupies space.
• Classification of Matter: Matter can be classified based on its physical state and chemical composition.
  • 2.1. Physical Classification (States of Matter):
    • Solid: Particles are tightly packed in fixed positions, giving solids a definite shape and volume. Intermolecular forces are very strong. They are incompressible and have low kinetic energy. (Examples: Ice, wood, metals)
    • Liquid: Particles are closely packed but can move freely past each other. Liquids have a definite volume but take the shape of their container. Intermolecular forces are moderate. They are slightly compressible and can flow. (Examples: Water, oil, mercury)
    • Gas: Particles are far apart and move rapidly and randomly in all directions. Gases have indefinite shape and volume (they fill the entire container). Intermolecular forces are negligible. They are highly compressible and diffuse quickly. (Examples: Air, oxygen, hydrogen)
    • Plasma: Often called the “fourth state of matter.” It consists of superheated, ionized gas where atoms have lost or gained electrons, forming a mixture of electrons and positive ions. It is an excellent conductor of electricity and is affected by magnetic fields. Found in stars, lightning, and fluorescent lamps.
    • Bose-Einstein Condensate (BEC): A state of matter formed when a gas of bosons is cooled to temperatures very close to absolute zero (-273.15 °C or 0 K). At this point, the individual atoms lose their individual identities and behave as a single quantum mechanical wave. Superfluidity and superconductivity are properties observed in BECs.
  • 2.2. Chemical Classification:
    • A. Pure Substances: Substances that have a definite and constant chemical composition and distinct properties. They cannot be separated into simpler constituents by physical methods.
      • i. Elements: The simplest form of pure substance that cannot be broken down into simpler substances by ordinary chemical means (e.g., heating, light, electricity). Each element is composed of only one type of atom.
        • Examples: Hydrogen (H), Oxygen (O), Iron (Fe), Gold (Au).
        • Types: Metals (conductors, malleable, ductile), Non-metals (poor conductors, brittle), Metalloids (intermediate properties, e.g., Silicon, Germanium).
      • ii. Compounds: Pure substances formed when two or more different elements combine chemically in a fixed ratio by mass. The properties of a compound are entirely different from its constituent elements.
        • Examples: Water (H2O), Carbon Dioxide (CO2), Sodium Chloride (NaCl).
        • Characteristics:
          • Constituent elements are in a fixed proportion by mass.
          • Properties are different from constituent elements.
          • Cannot be separated into constituent elements by physical methods.
          • Formation involves energy changes (absorption or release).
    • B. Mixtures: Substances containing two or more pure substances (elements or compounds) combined in any ratio, where each substance retains its individual properties. They can be separated by physical methods.
      • i. Homogeneous Mixtures: Mixtures where the components are uniformly distributed throughout, and the composition is uniform at the macroscopic level. Components cannot be seen separately.
        • Examples: Salt solution, sugar solution, air, alloys (e.g., brass).
        • Also known as solutions.
      • ii. Heterogeneous Mixtures: Mixtures where the components are not uniformly distributed, and the composition varies from one part to another. Components can often be seen separately.
        • Examples: Sand and water, oil and water, grain mixture, milk (colloidal dispersion).
        • Types: Suspensions, colloids.

3. Properties of Matter and Their Measurement

• 3.1. Physical Properties: Properties that can be measured or observed without changing the chemical identity of the substance.
  • Examples: Colour, odour, density, melting point, boiling point, hardness, conductivity.
• 3.2. Chemical Properties: Properties that describe how a substance reacts or changes its chemical composition when it interacts with other substances or energy.
  • Examples: Acidity, basicity, flammability, reactivity with water, oxidation states.
• 3.3. Measurement in Chemistry (SI Units):
  • International System of Units (SI): A coherent system of units adopted internationally.
  • Seven Base Units: meter | kg | second | ampere | Kelvin | mole | candela |
  • Derived Units: Units obtained by combining base units (e.g., density (kg/m3), volume (m3), speed (m/s)).
  • Mass and Weight: Mass is the amount of matter in a substance (constant). Weight is the force exerted by gravity on an object (varies with gravity).
  • Volume: Derived unit. 1 L = 1 dm3 = 1000 cm3 = 1000 mL.
  • Density: Mass per unit volume. Density = Mass / Volume. (Units: kg/m3 or g/cm3 or g/mL).
  • Temperature: Commonly measured in Celsius (°C), Fahrenheit (°F), and Kelvin (K).
    • K = °C + 273.15
    • °F = (9/5)°C + 32
    • 0 K is absolute zero (lowest possible temperature).

4. Uncertainty in Measurement

• 4.1. Precision and Accuracy:
  • Precision: Refers to the closeness of a set of measurements to each other. (Reproducibility).
  • Accuracy: Refers to the closeness of a single measurement (or average of measurements) to the true value.
• 4.2. Significant Figures:
  • Meaningful digits in a measured or calculated quantity that are known with certainty plus one uncertain digit. They reflect the precision of the measurement.
  • Rules for Determining Significant Figures:
    1. All non-zero digits are significant. (e.g., 285 cm has 3 sig figs).
    2. Zeros preceding the first non-zero digit are NOT significant. (e.g., 0.005 mL has 1 sig fig).
    3. Zeros between two non-zero digits ARE significant. (e.g., 2.005 g has 4 sig figs).
    4. Zeros at the end of a number ARE significant IF they are to the right of the decimal point. (e.g., 0.200 g has 3 sig figs, 100.0 has 4 sig figs).
    5. Zeros at the end of a number (without a decimal point) MAY or MAY NOT be significant. (e.g., 100 kg could have 1, 2, or 3 sig figs. Ambiguity is removed by using scientific notation: 1 x 10^2 (1 sf), 1.0 x 10^2 (2 sf), 1.00 x 10^2 (3 sf)).
    6. Exact numbers (e.g., number of objects, conversion factors like 1 dozen = 12) have infinite significant figures.
  • Rules for Arithmetic Operations with Significant Figures:
    1. Addition and Subtraction: The result must be reported with the same number of decimal places as the measurement with the fewest decimal places.
       • Example: 12.11 + 18.0 + 1.012 = 31.122 -> Round to 31.1 (1 decimal place from 18.0).
    2. Multiplication and Division: The result must be reported with the same number of significant figures as the measurement with the fewest significant figures.
       • Example: 2.5 x 1.25 = 3.125 -> Round to 3.1 (2 sig figs from 2.5).
• 4.3. Scientific Notation (Exponential Notation):
  • Used to represent very large or very small numbers concisely.
  • Numbers are expressed in the form N×10n, where N is a number between 1 and 10, and ‘n’ is an integer (positive or negative).
  • Example: 234,000 = 2.34×105; 0.00000016 = 1.6×10−7.
• 4.4. Dimensional Analysis (Factor Label Method):
  • A method for solving problems by using conversion factors (equalities between different units) to ensure that the units cancel correctly, leaving the desired units.
  • Example: Convert 2.5 cm to inches. (1 inch = 2.54 cm) 2.5 cm×(1 inch/2.54 cm)=0.98 inch

5. Laws of Chemical Combination

These fundamental laws govern how elements combine to form compounds.

• 5.1. Law of Conservation of Mass (Antoine Lavoisier):
  • Statement: Mass can neither be created nor destroyed in a chemical reaction.
  • Implication: The total mass of reactants must be equal to the total mass of products in a balanced chemical equation.
  • Example: When 12 g of Carbon reacts with 32 g of Oxygen, 44 g of Carbon Dioxide is formed. (12 + 32 = 44).
• 5.2. Law of Definite Proportions (Joseph Proust):
  • Statement: A given chemical compound always contains exactly the same proportion of elements by mass, irrespective of its source or method of preparation.
  • Example: Pure water (H2O) always contains Hydrogen and Oxygen in a 1:8 mass ratio, regardless of whether it’s river water or laboratory-prepared water.
• 5.3. Law of Multiple Proportions (John Dalton):
  • Statement: When two elements combine to form more than one compound, if the mass of one of the elements is fixed, then the masses of the other element bear a simple whole-number ratio to one another.
  • Example: Carbon and Oxygen form Carbon Monoxide (CO) and Carbon Dioxide (CO2).
    • In CO: 12 g C : 16 g O (Ratio of O to C = 16/12 = 1.33)
    • In CO2: 12 g C : 32 g O (Ratio of O to C = 32/12 = 2.66)
    • Keeping Carbon fixed at 12g, the masses of Oxygen are 16g and 32g. The ratio of these masses is 16:32 = 1:2, a simple whole-number ratio.
• 5.4. Gay-Lussac’s Law of Gaseous Volumes:
  • Statement: When gases combine or are produced in a chemical reaction, they do so in a simple whole-number ratio by volume, provided all reactants and products are at the same temperature and pressure.
  • Example: H2 (g) + Cl2 (g) -> 2HCl (g)
    • 1 volume of H2 reacts with 1 volume of Cl2 to give 2 volumes of HCl. (Ratio = 1:1:2).
• 5.5. Avogadro’s Law:
  • Statement: Equal volumes of all gases contain an equal number of moles (or molecules) under the same conditions of temperature and pressure.
  • Implication: At STP (Standard Temperature and Pressure: 0°C or 273.15 K, 1 atm or 101.325 kPa), 1 mole of any ideal gas occupies 22.4 liters.

6. Atomic Mass, Molecular Mass, and Mole Concept

• 6.1. Atomic Mass:
  • Atomic Mass Unit (amu) or Unified Mass (u): Defined as exactly one-twelfth (1/12) the mass of one atom of Carbon-12 isotope. 1 amu = 1.66056×10−24 g.
  • Relative Atomic Mass: The average relative mass of an atom as compared to 1/12 the mass of one Carbon-12 atom.
  • Average Atomic Mass: For elements having isotopes, it is the weighted average of the atomic masses of its isotopes based on their natural abundance.
    • Example: Chlorine (Cl) has two isotopes, Cl-35 (75%) and Cl-37 (25%). Average atomic mass = (35×0.75)+(37×0.25)=35.5 u.
• 6.2. Molecular Mass:
  • The sum of the atomic masses of all the atoms in a molecule.
  • Example: Molecular mass of H2O = (2 x Atomic mass of H) + (1 x Atomic mass of O) = (2 x 1.008 u) + (1 x 16.00 u) = 18.016 u.
• 6.3. Formula Mass:
  • Used for ionic compounds, which do not exist as discrete molecules but as a 3D network. It is the sum of atomic masses of all atoms in the empirical formula.
  • Example: Formula mass of NaCl = Atomic mass of Na + Atomic mass of Cl = 23 u + 35.5 u = 58.5 u.
• 6.4. Mole Concept:
  • Mole: The amount of substance that contains as many elementary entities (atoms, molecules, ions, electrons, etc.) as there are atoms in exactly 12 grams of the Carbon-12 isotope.
  • Avogadro’s Number (NA): The number of entities in one mole. NA=6.022×1023 entities/mol.
  • Molar Mass: The mass of one mole of a substance. Numerically equal to atomic mass, molecular mass, or formula mass in grams per mole (g/mol).
    • Example: Molar mass of H2O = 18.016 g/mol.
  • Relationships in Mole Concept:
    • Number of moles (n) = Mass of substance (g) / Molar mass (g/mol)
    • Number of moles (n) = Number of particles / Avogadro’s Number (NA)
    • Number of moles (n) = Volume of gas (L) / 22.4 L/mol (at STP)
• 6.5. Percentage Composition:
  • The percentage by mass of each element present in a compound.
  • Percentage of an element = (Mass of element in compound / Molar mass of compound) x 100.
• 6.6. Empirical Formula and Molecular Formula:
  • Empirical Formula: The simplest whole-number ratio of the atoms of different elements present in a compound.
  • Molecular Formula: The actual number of atoms of each element present in a molecule of the compound.
  • Relationship: Molecular Formula = (Empirical Formula)n, where n = (Molecular Mass / Empirical Formula Mass).
  • Steps to Determine Empirical and Molecular Formula:
    1. Convert mass % of each element to grams (assume 100 g sample).
    2. Convert grams to moles for each element.
    3. Divide mole values by the smallest mole value to get the simplest mole ratio.
    4. If not whole numbers, multiply by a suitable integer to get whole numbers. This gives the Empirical Formula.
    5. Calculate Empirical Formula Mass.
    6. Determine ‘n’ (n = Molecular Mass / Empirical Formula Mass).
    7. Multiply subscripts in Empirical Formula by ‘n’ to get Molecular Formula.

7. Stoichiometry and Stoichiometric Calculations

• Stoichiometry: The quantitative relationship between reactants and products in a balanced chemical equation. The coefficients in a balanced equation represent the number of moles of each substance.
• Limiting Reagent (Limiting Reactant): The reactant that is completely consumed during a chemical reaction, thereby limiting the amount of product that can be formed.
• Excess Reagent: The reactant that is not completely consumed and remains after the reaction is complete.
• Yields of Reaction:
  • Theoretical Yield: The maximum amount of product that can be formed from a given amount of reactants, as calculated from the balanced chemical equation (assuming 100% reaction efficiency).
  • Actual Yield: The amount of product actually obtained from a reaction, which is almost always less than the theoretical yield due to various factors (e.g., incomplete reaction, side reactions, loss during purification).
  • Percentage Yield: Percentage Yield = (Actual Yield / Theoretical Yield) x 100.

8. Concentration Terms

Various ways to express the concentration of a solution.

• 8.1. Mass Percentage (% w/w):
  • Mass % = (Mass of solute / Mass of solution) x 100
• 8.2. Volume Percentage (% v/v):
  • Volume % = (Volume of solute / Volume of solution) x 100
• 8.3. Mass by Volume Percentage (% w/v): (Common in pharmaceuticals)
  • Mass by Volume % = (Mass of solute (g) / Volume of solution (mL)) x 100
• 8.4. Parts Per Million (ppm) / Parts Per Billion (ppb):
  • Used for very dilute solutions.
  • ppm = (Mass of solute / Mass of solution) x 10^6
  • ppb = (Mass of solute / Mass of solution) x 10^9
• 8.5. Mole Fraction (X):
  • The ratio of the number of moles of one component to the total number of moles of all components in the solution.
  • X(A) = n(A) / (n(A) + n(B) + …)
  • Sum of mole fractions of all components in a solution is always 1. (X(A) + X(B) = 1)
• 8.6. Molarity (M):
  • Number of moles of solute dissolved per liter of solution.
  • M = Moles of solute / Volume of solution (L)
  • Units: mol/L or M.
  • Temperature Dependent: Molarity changes with temperature because volume changes with temperature.
  • Dilution Formula: M1V1 = M2V2
• 8.7. Molality (m):
  • Number of moles of solute dissolved per kilogram of solvent.
  • m = Moles of solute / Mass of solvent (kg)
  • Units: mol/kg or m.
  • Temperature Independent: Molality does not change with temperature as mass does not change with temperature.
• 8.8. Normality (N):
  • Number of gram equivalents of solute dissolved per liter of solution.
  • N = Gram equivalents of solute / Volume of solution (L)
  • Gram equivalents = Mass of solute / Equivalent mass.
  • Equivalent mass = Molar mass / Valency factor (n-factor).
  • Relationship between Molarity and Normality: N = M x Valency factor (n-factor)
  • Units: Eq/L or N.
  • Temperature Dependent.