Atomic and Molecular Masses

Think of relative atomic mass (Ar) as a comparison system - it tells you how heavy an atom is compared to a carbon-12 atom as the reference point. Scientists use 1/12th of a carbon-12 atom's mass as the standard unit called amu (atomic mass units).

The clever bit is that relative atomic mass accounts for isotopes - different versions of the same element with varying masses. It calculates an average based on how common each isotope is in nature.

Relative molecular mass (Mr) works the same way but for entire molecules. You simply add up all the Ar values of the atoms in the compound to get the total molecular mass.

Quick Tip: Always check your periodic table for Ar values - they're the decimal numbers under each element symbol!

The Mole and Gas Laws

The mole is chemistry's counting unit - like saying "a dozen" but for atoms and molecules. One mole always contains 6.02 × 10²³ particles (that's Avogadro's constant), whether you're counting atoms, ions, or molecules.

This massive number lets you convert between the microscopic world of atoms and the macroscopic world you can measure. The formula Number of particles = moles × Avogadro's constant is your key tool here.

For gases specifically, the ideal gas equation pV = nRT connects pressure, volume, temperature, and moles. Remember your units: pressure in Pa, volume in m³, temperature in Kelvin $°C + 273$, and R = 8.31 J mol⁻¹ K⁻¹.

Concentration measures how much substance is dissolved per unit volume, typically in mol/dm³. This stays constant throughout a solution, making it perfect for calculations.

Memory Aid: Think of moles as chemistry's "packaging system" - it groups particles into manageable, countable units!