Quick questions on Dilution, concentration units and pH on dilution explained: HSC Chemistry Module 6

Molarity is the default HSC unit because it links directly to stoichiometry ($n = cV$). The other units appear in environmental and biological contexts (lead in drinking water, residual chlorine, salinity).

When solvent is added to a solution, the moles of solute do not change. Therefore:

For a strong acid, $[H^+] = c$ (fully ionised). On dilution, $[H^+]$ drops in direct proportion to $c$. A 10-fold dilution raises pH by 1.0 unit; a 100-fold dilution raises pH by 2.0 units.

If you need a very dilute solution, a single one-step dilution can be inaccurate (pipetting a very small volume). Serial dilution does it in stages: each step is a 10- or 100-fold dilution. To go from 1.00 mol/L to $1.00 \times 10^{-4}$ mol/L, do four successive 10-fold dilutions.

A 10-fold dilution raises pH by:

$[H^+] = 0.0100$ mol/L, $pH = 2.00$.

$c = 1.00 \times 10^{-4}$ mol/L, $pH = 4.00$.

$c = 1.00 \times 10^{-4}$ mol/L:

In $c_1 V_1 = c_2 V_2$, both volumes must be in the same unit. Mixing mL and L drops a factor of 1000.

That formula is for strong acids only. Use $\sqrt{K_a c}$ for weak acids.

$10^{-9}$ mol/L $HCl$ does not have pH 9. Below about $10^{-6}$ mol/L, water's auto-ionisation contributes; a full charge balance gives pH approaching 7 from the acidic side.

Dilution lowers concentration; it does not change the strength (degree of ionisation). A dilute strong acid is still a strong acid (still fully ionised).

A 5 percent w/v $NaOH$ solution is 50 g/L, which converts to $50/40 = 1.25$ mol/L. Always carry the conversion through molar mass.

When making a standard, the pipette should be rinsed with the stock first (not water), and the volumetric flask should be rinsed with distilled water (not stock).