Explicitly identifying chemical amount as a number and absolute temperature as an energy focuses student efforts on mastering these broadly applicable concepts and demonstrates that chemistry is the foundational science it claims to be rather than an arcane specialty.  The equivocation in general chemistry texts between the SI definition of the mole as the dimensioned unit of chemical amount and the simpler definition of the mole as a dimensionless unit of number sows seeds of confusion well beyond stoichiometry.  The mole should be unambiguously defined as a dimensionless unit, specifically the number of atomic mass units in a gram (1mol = 1g/amu = 6.02..x10²³), and clearly distinguished from its unit conversion factor (N_A = 6.02..x10²³mol^-1 = 1).  Formula masses are then microscopic masses per formula unit, conveniently expressed as the macroscopic mass of a large number of formula units divided by Avogadro's large number, reaction energies are microscopic energy changes per reaction, conveniently expressed as macroscopic energies divided by one mole, and Faraday's constant is the absolute charge of the electron (F = 9.65..x10⁴C∙mol^-1 = 1.60..x10^-19C = e).  Similarly, general chemistry texts miss an opportunity to ground absolute temperature in the atomic nature of matter.   Absolute temperature should be defined as the PV energy per molecule in a dilute gas thermometer (T = PV/n as n/V→0).  The gas constant and Boltzmann's constant are then unit conversion factors between the kelvin unit of energy and other units of energy (R = 8.31..J∙mol^-1∙K^-1 = 1 = 1.38..x10^-23J∙mol^-1 = k_B), formula heat capacities are dimensionless ratios, and reaction entropies are changes in the natural logarithm of the number of possible arrangements per reaction.