Introduction
The microscopic world of particles produces problems of scale in terms of their almost infintesimally small size and the enormous numbers of them.
This problem was dealt with specifically by Amedeo Avogadro, who developed the work of French chemist Joseph Louis Guy-Lussac (1778-1850) with the 1811 publication of his hypothesis, and the idea that gases are made up of atoms or combinations of atoms (molecules) and can be quantified. Although his work was largely ignored during his lifetime, by the 1880s it was universally accepted, thanks to Stanislao Cannizzaro, who created a table of atomic weights based on Avogadro's work.
This section deals with this quantification and its utility to modern chemists.
Structure 1.4.1 - The mole (mol) is the SI unit of amount of substance. One mole contains exactly the number of elementary entities given by the Avogadro constant.
- Convert the amount of substance, n, to the number of specified elementary entities.
Guidance
- An elementary entity may be an atom, a molecule, an ion, an electron, any other particle or a specified group of particles.
- The Avogadro constant NA is given in the data booklet. It has the units mol–1.
Tools and links
Structure 1.4.2 - Masses of atoms are compared on a scale relative to 12C and are expressed as relative atomic mass Ar and relative formula mass Mr .
- Determine relative formula masses Mr from relative atomic masses Ar .
Guidance
- Relative atomic mass and relative formula mass have no units.
- The values of relative atomic masses given to two decimal places in the data booklet should be used in calculations.
Tools and links
- Structure 3.1 - Atoms increase in mass as their position descends in the periodic table. What properties might be related to this trend?
Structure 1.4.3 - Molar mass M has the units g mol–1.
- Solve problems involving the relationships between the number of particles, the amount of substance in moles and the mass in grams.
Guidance
- The relationship n = m/M is given in the data booklet.
Tools and links
- Reactivity 2.1 - How can molar masses be used with chemical equations to determine the masses of the products of a reaction?
Structure 1.4.4 - The empirical formula of a compound gives the simplest ratio of atoms of each element present in that compound. The molecular formula gives the actual number of atoms of each element present in a molecule.
- Interconvert the percentage composition by mass and the empirical formula.
- Determine the molecular formula of a compound from its empirical formula and molar mass.
Guidance
Tools and links
- Tool 1 - How can experimental data on mass changes in combustion reactions be used to derive empirical formulas?
- Nature of science, Tool 3, Structure 3.2 - What is the importance of approximation in the determination of an empirical formula?
Structure 1.4.5 - The molar concentration is determined by the amount of solute and the volume of solution.
- Solve problems involving the molar concentration, amount of solute and volume of solution.
Guidance
- The use of square brackets to represent molar concentration is required.
- Units of concentration should include g dm–3 and mol dm–3 and conversion between these.
- The relationship n = CV is given in the data booklet.
Tools and links
- Tool 1 - What are the considerations in the choice of glassware used in preparing a standard solution and a serial dilution?
- Tool 1, Inquiry 2 - How can a calibration curve be used to determine the concentration of a solution?
Structure 1.4.6 - Avogadro’s law states that equal volumes of all gases measured under the same conditions of temperature and pressure contain equal numbers of molecules.
- Solve problems involving the mole ratio of reactants and/or products and the volume of gases.
Guidance
Tools and links
- Structure 1.5 - Avogadro’s law applies to ideal gases.
- Under what conditions might the behaviour of a real gas deviate most from an ideal gas?
Structure 1.4 - Counting particles by mass: The mole