History of The Periodic Table

This area of study focuses on the historical development of, and the relationship between, the Periodic Table and atomic theory. Students investigate trends and patterns within the Periodic Table and use subshell notation to describe the electronic configuration of elements. They explore the link between the electronic configuration of an element and the type of bonding in which it participates. Students are introduced to many of the major qualitative and quantitative ideas fundamental to chemistry including empirical and molecular formulas and the mole concept. They undertake practical activities that build their understanding of the Periodic Table.

Stoichiometry

Particles and the Mole

Particles

Matter: anything that possesses mass, and takes up space.

Particle: molecules, electrons, protons, neutrons and the like.

Atom: the smallest piece of a chemical element that retains its unique chemical properties.

Molecule: at least two atoms, held together by chemical bonds.

You may remember that the matter of which the universe is composed is itself made up of particles. Chemistry deals with many particles: all molecules and the atoms that make them up are specific types of particle, we may also need to understand the behaviour of smaller subatomic particles, such as the electron or proton.

In general, the particles which we will encounter will all be of minuscule size and mass, and in vast numbers - a single carbon atom weighs and there are more than 5×10²² of them in a single gram of graphite.

The Mole

Problems with Size

It can become very cumbersome to deal with such large and small numbers, especially during calculations. To overcome this impracticality, chemists deal with multiples of particles, instead of individually. A chemist will avoid calculating directly the number particles in a given sample - if the sample is of everyday proportions and the particles are on a molecular scale, then the number of particles will be immense whilst the mass of each particle will be infinitesimal. If instead the chemist calculated the number of multiples of a large, fixed number of particles, then, with clever selection of the specific number, the number of multiples present and the mass of each multiple would be much simpler and managable numbers. This can be compared to the sale of tiny items: if for instance, a grain merchant was to sell kernels of grain individually, then the price of each kernel would be frustratingly small and the merchant would need to painstakingly count out each kernel in order to make a sale. It is much more practical for the merchant to sell the kernels in intervals of some enormous number (by the kilogram, or even the ton) as they would be both easier to price and measure.

Atomic and Molecular Mass, the Mole

Theoretically, scientists could have chosen any large number and used it as a standard multiple. It would seem convenient to select a power of ten, like 10²³ or 10²⁴ as this would simplify calculations. Instead, scientists have chosen the mole, which is defined as approximately 6.02252×10²³ particles of the specific type in question.

Why 6.02×10²³? From first glance, this additional number seems to complicate rather than simplify matters. This number was however selected carefully and deliberately. The secret to its advantage lies in the mass of the particles under scrutiny.

Recall that each atom of a single chemical element has a specific atomic mass that is unique to that element, and measured in atomic mass units (AMU). Each AMU is defined as one twelfth the mass of a specific isotope of carbon, but an AMU can also be expressed in grams - it is about

Spectroscopy

AAS (atomic absorption spectroscopy) UV light and visible wavelengths are analysed. Works only with metals (most cations, some anions). Results obtained are qualitative.

VCE Chemistry/Printable version

Contents