Molecular Mass: Collective Atomic Weights in Molecules

• Definition: Sum of Atomic Weights in Molecules: The molecular mass of a substance is the sum of the atomic masses of all the atoms in a molecule of the substance. 
• It is therefore the relative mass of a molecule expressed in atomic mass units (u).
  • Example: The molecular mass of HNO3 = the atomic mass of H + the atomic mass of N + 3 × the atomic mass of O

= 1 + 14 + 3×16 = 63 u

Calculating the Ion Contributions in Compounds:

• Understanding Atomic Weights in Compounds: The formula unit mass of a substance is a sum of the atomic masses of all atoms in a formula unit of a compound.
• Calculation: Bridging Molecular Mass Techniques for Ionic Compounds: It is calculated in the same manner as we calculate the molecular mass. 
  • The only difference is that the word formula unit is used for those substances whose constituent particles are ions. 
  • Example: Formula unit mass of sodium chloride (NaCl) can be calculated as:

1 × 23 + 1 × 35.5 = 58.5 u

What role does the Mole concept play in Chemistry, particularly in relation to Molecular Mass?

• Development of Mole Concept: From ‘Heap’ to Fundamental Unit in Chemistry: The word “mole” was introduced around 1896 by Wilhelm Ostwald who derived the term from the Latin word moles meaning a ‘heap’ or ‘pile’. 
  • A substance may be considered as a heap of atoms or molecules. 
  • In 1967, the unit mole was accepted to provide a simple way of reporting a large number.
• Understanding Concept: Bridging Quantities and Mass in Chemistry: One mole of any species (atoms, molecules, ions or particles) is that quantity in number having a mass equal to its atomic or molecular mass in grams.
  • The mole, symbol mol, is the SI unit of amount of substance. 
• Avogadro Constant: Magnitude of Mole-Particle Relationships in Chemistry The number of particles (atoms, molecules or ions) present in 1 mole of any substance is fixed, with a value of 6.02214076 x 10²³.  
  • This number is called the Avogadro Constant or Avogadro Number (represented by N₀), named in honor of the Italian scientist, Amedeo Avogadro.
• 1 mole (of anything) = 6.02214076 x 10²³ in number  = Relative mass in grams
• Relativity in Chemistry: Connecting Mole and Molecular Mass: The mass of 1 mole of a substance is equal to its relative atomic or molecular mass in grams. 
• Atomic Mass: Weight of Individual Atoms: The atomic mass of an element gives us the mass of one atom of that element in atomic mass units (u). 
• Gram Atomic Mass: Link Between Atomic and Molar Mass in Chemistry: Mass of 1 mole of an atom of that element, that is, molar mass is also known as gram atomic mass. 
  • Example: Atomic mass of oxygen =16u. So, gram atomic mass of oxygen = 16 g.
    • 16 u oxygen has only 1 atom of oxygen, 16 g oxygen has 1 mole atom, that is, 6.022x10²³ atoms of oxygen.