CHEM1201 2.2- Chemical Formulae, Molar Mass and Balancing Equations

Page 1: Molar Mass Calculations

  • Overview of Molar Mass Calculation Steps

    1. Determine the number of atoms of each type in the compound.

    2. Find the atomic mass of each atom in the compound.

    3. Calculate the sum of the atomic masses to find the molar mass.

  • Example: Iron(III) Oxide (Fe2O3)

    1. Atoms: 2 x Fe and 3 x O

    2. Atomic Masses: Fe = 56 g/mol; O = 16 g/mol

    3. Molar Mass: (2 x 56) + (3 x 16) = 160 g/mol

  • Additional Molar Mass Calculations

    • a) Fe2O3

      • Atomic masses: Fe = 56 g/mol; O = 16 g/mol

      • Molar mass = (2 x 56) + (3 x 16) = 160 g/mol

    • b) Cl2

      • Atomic mass: Cl = 35.5 g/mol

      • Molar mass = (2 x 35.5) = 71 g/mol

    • c) Al2(SO4)3

      • Atomic masses: Al = 27 g/mol; S = 32 g/mol; O = 16 g/mol

      • Atoms: 2 x Al, 3 x S, 12 x O

      • Molar mass = (2 x 27) + (3 x 32) + (12 x 16) = 342 g/mol

  • Mass from Moles Calculations

    • a) Fe2O3

      • mass = n x molar mass = 0.25 mol x 160 g/mol = 40 g

    • b) Cl2

      • mass = n x molar mass = 0.25 mol x 71 g/mol = 17.8 g

    • c) Al2(SO4)3

      • mass = n x molar mass = 0.25 mol x 342 g/mol = 85.5 g

Page 2: Relating Moles and Number of Atoms

  • Avogadro’s Number

    • 1 mole contains 6.02 x 10^23 atoms.

  • Atoms in Given Moles

    • a) 0.25 mole of Fe

      • number of atoms = 0.25 mole x 6.02 x 10^23 atoms/mole = 1.51 x 10^23 atoms

    • b) 0.25 mole of H2O

      • One molecule has 3 atoms (2 H and 1 O).

      • Atoms = 0.25 x 3 x 6.02 x 10^23 = 4.52 x 10^23 atoms

  • Summary of Key Concepts

    • Moles = measurement of substance amount

    • Avogadro’s number (NA) connects moles to number of molecules

    • Molar mass is derived from chemical formulae and atomic mass

    • Understand mass relationship with moles and molar mass.

    • Note: Valency and periodic table position are important.

  • Elemental Formulae

    • Single atoms (e.g., Sodium Na) vs diatomic molecules (e.g., Cl2).

Page 3: Periodic Table and Diatomic Molecules

  • Diatomic Elements

    • List of Elements forming Diatomic Molecules:

      • H2, N2, O2, F2, Cl2, Br2, I2

      • Notable exceptions: S8 (sulfur) and P4 (phosphorus)

  • Importance of Knowing Diatomic Elements

    • Knowing these ensures accurate representation in equations, which is crucial for chemical reactions.

Page 4: Recognizing Chemical Reactions

  • Signs of a Chemical Reaction

    • Changes such as color change, gas production, and temperature changes.

    • Emission or absorption of light is also an indicator.

  • Chemical Reactions

    • Involves reactants on the left, products on the right.

    • Example:

      • Sodium and Chlorine react to form Sodium Chloride (Na + Cl2 -> NaCl)

  • Steps to Describe a Chemical Reaction

    1. Identify what happens.

    2. Write correct formulae for each species.

    3. Balance the chemical equation to ensure equal atom numbers.

Page 5: Balancing Combustion Reactions

  • Balancing Reactions: Example with Methane (CH4)

    • Reactions involve breaking and forming chemical bonds.

    • Balance oxygen and hydrogen: CH4 + 2O2 -> CO2 + 2H2O

    • One mole of CH4 reacts with two moles of O2.

  • Example of Ammonia Combustion Reaction:

    • Reaction: 2NH3 + 3O2 -> 2N2 + 6H2O

    • Ratios remain constant across different formats.

Page 6: Conclusions on Chemical Reaction Principles

  • Mass Conservation

    • Total number of atoms remains unchanged; mass of reactants = mass of products

    • Example: 58g of butane with 208g of oxygen gives 176g of CO2 + 90g of H2O.

  • Balancing Chemical Equations

    • Only coefficient numbers can be changed to balance equations, subscripts must not be altered.

    • Example: CO + O2 -> CO2 (2CO + O2 -> 2CO2)

  • Practice Problems

    • Balance equations: 1) Ca + O2 -> CaO

      1. KClO3 -> KCl + O2

      1. Al2O3 + H2SO4 -> Al2(SO4)3 + H2O