Elements, compounds and mixtures Study Guide

Dmitri Mendeleev's initial Periodic Table, developed in 1869, was a revolutionary step in chemistry. He arranged elements primarily by increasing atomic mass. He noticed recurring patterns in chemical properties and used these patterns to group elements into columns (groups). Crucially, Mendeleev understood the importance of these periodic properties so much that he sometimes deliberately left gaps in his table and even reversed the order of some elements if their properties suggested they belonged elsewhere, despite their atomic masses. For instance, tellurium (atomic mass 127.6) was placed before iodine (atomic mass 126.9) because iodine's properties clearly aligned with the halogens, and tellurium's with the oxygen group. These 'anomalies' were later explained with the discovery of subatomic particles and the development of the concept of atomic number.

The refinement of Mendeleev's table into the modern Periodic Table came with the understanding that it is the number of protons in an atom's nucleus, known as the atomic number (Z), that determines its chemical identity and properties. This understanding emerged with the work of Henry Moseley in the early 20th century. Moseley used X-ray spectroscopy to determine the atomic number of elements, confirming that each element has a unique number of protons. When elements are arranged strictly by increasing atomic number, the few discrepancies (like tellurium and iodine) in Mendeleev's original mass-based ordering are resolved, and the periodic recurrence of chemical properties becomes even more consistent. The modern Periodic Table, therefore, organizes elements in order of increasing atomic number, which naturally aligns elements with similar electron configurations in the same groups, explaining their similar chemical behaviours.

When a substance changes state, such as melting or boiling, energy is transferred to or from it. This energy does not increase the kinetic energy of the particles (and therefore not the temperature) during the change of state; instead, it is used to overcome the forces holding the particles together. In solids, particles are held in fixed positions by strong forces of attraction (chemical bonds or intermolecular forces). To melt a solid, enough energy must be supplied to partially overcome these forces, allowing particles to move past each other, forming a liquid. The melting point is the temperature at which this occurs. For boiling, even more energy is needed to completely overcome the forces, allowing particles to escape as a gas. The boiling point is the temperature at which this happens. The stronger the forces between particles, the more energy is required to overcome them, resulting in higher melting and boiling points. For example, substances with strong covalent bonds (like diamond) or ionic bonds (like sodium chloride) have very high melting and boiling points because a large amount of energy is needed to break these strong bonds. Substances with weaker intermolecular forces, such as simple molecular substances (like water or oxygen), require less energy to change state and therefore have lower melting and boiling points. The amount of energy transferred relates directly to the strength of these forces, explaining the different temperatures observed for changes of state.