Atomic Model and the Periodic Table
The history behind the atomic model up to Bohr, and also the trends in the periodic table...
History of the atomic model
In 1803, John Dalton proposed was made of solid indivisible particles which he called atoms. He reasoned that the different elements had different atomic weights.
In 1897, J.J Thomson discovered the electron, which led to a new model being developed for the atom, referred to as the plum pudding model. This new model proposed that the atom was a positive sphere in which the negatively charged electrons were embedded into.
In 1911, Ernest Rutherford developed the planetary model of the atom, following his discovery of the nucleus in the centre of the atom. This model proposed that the negatively charged electrons orbit a central positively charged nucleus, similar to the way the planets orbit the sun.
In 1913, Niels Bohr inferred correctly that the electrons must occupy stable, non-radiating orbits called shells (or energy levels) around the central positively charged nucleus. In addition to this, Bohr released that electrons could move from one energy level to another, through the absorption or emission of a fixed amount of energy.
In 1927-35, Erwin Schrodinger refined Bohr’s model into the cloud model of the atom. The main difference between Bohr’s model and Schrodinger’s model, was that unlike Bohr, Schrodinger proposed orbitals (or clouds) which are regions of space where electrons have a higher probability of being located, rather than the precise orbits which Bohr’s model suggests.
Periodic table development
In 1789, Antoine Lavoisier classified elements into four sub-categories: metals, non-metals, elastic fluids (gases) and earths, some of which (the earths) were shown later to be compounds. Lavoisier also supported the idea of accurate gravimetric analysis and the conservation of mass in a chemical reaction.
In 1817-1829, Johann Doberener observed notable relationships between groups of elements, which he called triads. He noted that not only the properties of the middle element were between the other two elements, but also had an atomic weight that was close to the average of the other two.
In 1864, John Newlands, following Doberener’s work arranged the elements into a table with increasing atomic weight. He noted that every eighth element after a given element possessed similar physical and chemical properties. However, his table broke down, as it paired non-metals such as nitrogen and phosphorus with metals such as manganese.
In 1869-71, Dmitri Mendeleev compared the similarities in chemical and physical properties to produce his own classification table, which was based on atomic weights. Where the atomic weight of the element was not accurately known, he placed the element in a position consistent with its properties. He left spaces for undiscovered elements in the belief that the properties of these undiscovered elements could be predicted by analysing trends across the rows and columns of the table. Elements with similar properties occupied vertical groups.
In 1870, Lothar Meyer developed a periodic table similar to that of Mendeleev. However because Mendeleev’s table is based on chemical properties, generalisations and predictions could be made more readily.
In 1893-1898, William Ramsay along with Lord Rayleigh discovered the unreactive noble gases. Ramsay though that this new family should be added to the right-hand end of Mendeleev’s table.
In 1913, Anton van den Broek suggested that the elements of the periodic table be arranged according to the charge on their nucleus rather than according to their atomic weight.
In 1912-13, Henry Moseley conducted a series of experiments to investigate Broek’s ideas, where he measured the X-ray spectra of the first ten elements and noted a fundamental quantity for each element, which is referred to as atomic number. As a result, the current periodic table is organised by atomic number. Periodic law: The properties of elements vary periodically with their atomic numbers.
Periodic table trends Electrical conductivity
Electrical conductivity is related to the metallic nature of the element. The more metallic the element, the better the conductivity. As a result metals are good conductors the non-metals are poor conductors.
This is because metals form positive ions, resulting in free electrons that are able to flow through the substance. Due to these free electrons allowing electric current to flow.
Generally, conductivity decreases across a period, and increases downwards towards the period.
Atomic radius is the size of the atom and is dependent on the number of shells and the magnitude of charge of the nucleus.
Atomic radius decreases across the period, from left to right. This is because the increased positive charge pulls the outermost electrons (same energy level) closer to the nucleus.
Atomic radius increases down the group, from top to bottom. This is because there are additional shells.
Ionisation energy is the energy required to remove an electron from an atom of an element in a gaseous state.
The first ionisation energy is the energy required to remove the first electron
Second ionisation energy is the energy required to remove the second electron. It is always greater than the first ionisation energy as it requires more energy to remove an electron from a positive ion due to the extra electrostatic attraction.
Third ionisation energy is the energy required to remove the third electron, and so on. The third ionisation is always greater than the second ionisation energy as it is even more positively charged.
Ionisation energies provide strong evidence for periodic law. This is because ionisation energy is a direct result of the increasing atomic number.
The first ionisation energy increases across each period with a minimum at the alkali metals (Group I) and a maximum at the noble gases (Group VIII). This is due to each successive element have one additional proton resulting in an increase of attractive force.’
For second ionisation energy, the minimum is at the alkaline earth metals (Group II) and increases to the noble gases (Group VIII), with the maximum being the alkali metals (Group I) and so on.
Ionisation energy decreases down the group, as outer electrons are further away from the nucleus, as the atomic radius increases.
With Group I and Group II metals, the reactivity of the metal is related to the first ionisation energy. With transitions metals the relationship is not as simple.
Melting points and boiling points
Melting point and boiling points are related to the strength of bonding within the substance. The stronger the bonding within the substance, the higher the melting points and boiling points.
Metallic bonding strength varies, but is strongest in the transition metals. Metals generally have a high melting point, and a high boiling point.
Covalent molecular forces are very weak, so covalent molecular substances have a low melting and boiling points, but covalent network bonding is very strong, so they have high melting and boiling points.
Melting point and boiling points generally rise up to Group IV (carbon, silicon and germanium) and then decreases.
The trend down a group is variable, because it is dependent on the bonding and type of crystal lattice.
It generally decreases for groups I to IV (i.e. up to carbon)
It generally increases for groups V to VIII, and also with the transition metals.