Analytical Chemistry: Volumetric Analysis Lecture 2

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Study Notes

Volumetric analysis involves the calculation of mole fractions of solutes in a solution.
A weak acid is a substance that does not completely dissociate in water, such as HF, HClO3, and HF.
The mole fraction of a solute can be calculated using the formula: mole fraction = number of moles of solute / total number of moles of solution.

Concentration units include:
- g/L (grams per liter)
- ppm (parts per million)
- ppb (parts per billion)
- w/w% (weight/weight percentage)
- w/v% (weight/volume percentage)
Conversion between concentration units can be done using the following formulas:
- ppm = (mass of solute / mass of solution) x 10^6
- ppb = (mass of solute / mass of solution) x 10^9

Arrhenius definition:
- Acids are H+ donors
- Bases are OH- donors
Broadened definition:
- Acids increase H+ concentration or [H+] increases
- Bases increase H+ acceptor
Arrhenius theory:
- Acids are substances that increase the concentration of the hydronium ion (H3O+) when dissolved in water
- Bases are substances that dissociate into hydroxide ions (OH-) and cations in solution
Problems with Arrhenius theory:
- H3O+ is not present in solution, but rather OH(H2O)3-
- Other substances also have acidic or basic properties

A Bronsted-Lowry acid is a species that donates a proton (H+)
A Bronsted-Lowry base is a substance that can accept a proton (H+)
Bronsted-Lowry acid-base theory has several advantages over Arrhenius theory:
- It can explain the behavior of substances that do not contain OH- or H3O+
- It can explain the behavior of substances that have multiple acidic or basic properties

A Lewis acid is a chemical compound that has a tendency to accept an electron pair or more from a donor compound
A Lewis base is a chemical compound that has a tendency to donate an electron pair to an acceptor compound
Examples of Lewis acids and bases include trimethylborane (Me3B) and ammonia (NH3)

The level of acidity is governed by acid strength, which is the result of the dissociation of acids in solvents
If the acid strength is greater than the solvent strength, the equilibrium will lie towards the right (100% dissociation of the acid)

Acid strength is defined by the equilibrium position of its dissociation reaction
A strong acid is one for which the equilibrium lies far to the right (almost all of the original acid is dissociated)
A weak acid is one for which the equilibrium lies far to the left (only a small amount of the original acid is dissociated)
Examples of strong acids include hydrochloric acid (HCl) and sulfuric acid (H2SO4)
Examples of weak acids include phosphoric acid (H3PO4), acetic acid (CH3COOH), and benzoic acid (C6H5COOH)

pH can be calculated using the formula: pH = -log[H3O+] = -log[H+]
The pH of a strong acid solution can be calculated using the concentration of the acid
The pH of a weak acid solution can be calculated using the equilibrium constant (Ka) and the concentration of the acid