SSC Chemistry Chapter 3: Atomic & Molecular Mass

Chapter 3: Basics of Matter Structure

• This chapter, the 3rd of SSC Chemistry, is about the structure of matter.
• The video series covers the chapter topic by topic.

Atomic Mass and Molecular Mass

• Dimitri Mendeleev published the periodic table in 1869.
• Elements are arranged by increasing atomic number in the periodic table.
• Each element has a unique atomic number which defines its position.
• Hydrogen's atomic number is 1, making it the first element.
• Atomic mass is the mass of an atom, measured in Atomic Mass Units (AMU).
• Molecular mass is the sum of the atomic masses in a molecule.
• Isotopes are atoms of the same element with different neutron numbers.
• Hydrogen isotopes include Protium, Deuterium (one neutron), and Tritium (two neutrons).

Calculating Atomic Mass

• Determining atomic mass involves: Σ (isotope mass x relative abundance).
• Determining relative atomic mass of chlorine uses the masses and abundances of its isotopes Cl-35 and Cl-37 as an example.
• Cl-35 has an abundance of 75.77%, while Cl-37 has 24.23%.

Gram Atomic Mass and Molar Mass

• Gram atomic mass is atomic mass expressed in grams (g/mol).
• Gram molecular mass, or molar mass, is molecular mass in grams (g/mol).
• One mole of any substance contains Avogadro's number of particles (6.02 x 10^23).
• Molar mass examples:
  • Oxygen (O2): 32 g/mol
  • Hydrogen (H2): 2 g/mol
  • Water (H2O): 18 g/mol
  • Carbon Dioxide (CO2): 44 g/mol

Volume of Gases

• At standard conditions (0°C and 1 atm), 1 mole of any gas occupies 22.4 liters (molar volume).
• Standard temperature is 0°C, and standard pressure is 1 atmosphere.
• Molar volume is applicable to gasses only.

Calculations using Moles

• To find the number of molecules in a given amount of water
  • One mole of water weighs 18 grams and contains 6.02 x 10^23 molecules.
• To calculate the mass of a single molecule of sulfuric acid (H2SO4)
  • H2SO4 has a molar mass of 98 g/mol.
  • One molecule's mass = 98 / (6.02 x 10^23) grams.
• To determine the volume of 5 moles of carbon dioxide (CO2) gas at standard conditions
  • One mole occupies 22.4 liters, thus 5 moles occupy 5 x 22.4 = 112 liters.

Concentration of Solutions

• Solutions are composed of a solute dissolved in a solvent.
• Molarity (M) measures concentration as moles of solute per liter of solution.
• Dilute solutions contain small solute amounts, whereas concentrated solutions contain large amounts.

Preparing Solutions

• A volumetric flask is often used in preparing a solution of specific molarity.
• Calculations involve the formula: M1V1 = M2V2, where M is molarity and V is volume.
• For example, to prepare 250 ml of 0.2 M solution from a 1 M stock solution, use 50 ml of the stock solution and dilute to 250 ml.

Percentage Composition

• Percentage composition indicates the relative mass of each element in a compound.
• To find it, divide the mass of each element by the compound's total mass, then multiply by 100.
• In H2SO4, determine the mass of H, S, and O, then divide by the total molar mass (98 g/mol).

Empirical and Molecular Formulas

• The empirical formula gives the simplest whole-number ratio of atoms in a compound.
• The molecular formula represents the actual number of atoms in a molecule.
• Steps to determine the empirical formula:
  • Convert percentage composition to grams.
  • Convert grams to moles.
  • Divide each mole value by the smallest mole value to get the simplest ratio.
• To find the molecular formula:
  • Calculate the empirical formula mass.
  • Determine n = (molecular mass) / (empirical formula mass).
  • Multiply the subscripts in the empirical formula by n to obtain the molecular formula.