Topic 8: The Periodic Table

The Periodic Table organises every known element into a logical grid based on the structure of their atoms. Once you understand the pattern behind the rows and columns, you can predict how an unfamiliar element will behave just from where it sits. This topic shows you how to read those patterns for the main groups and the transition block.

Arrangement by proton number

In the modern Periodic Table the elements are placed in order of increasing proton number (also called atomic number). Reading from left to right and top to bottom, each successive element has one more proton than the one before, so hydrogen (1) comes first and helium (2) next. This ordering is important because the number of protons fixes the number of electrons in a neutral atom, and electrons are what control chemical behaviour. Earlier chemists tried arranging by relative atomic mass, but a few pairs would have fallen out of place; using proton number puts every element exactly where its properties fit. The table is divided into metals on the left and centre, and non-metals on the upper right, separated by a rough diagonal staircase. Elements close to that boundary, such as silicon, can show some properties of both and are sometimes called metalloids.

Groups, periods and electronic configuration

The vertical columns are called groups and the horizontal rows are called periods. Elements in the same group have the same number of electrons in their outer shell, which is why they react in similar ways and form a chemical family. The group number (using the older numbering I to VIII) tells you the number of outer-shell electrons, so a Group I element has one outer electron and a Group VII element has seven. The period number tells you how many electron shells an atom has; an element in Period 3 has three occupied shells. For example, sodium has the configuration 2,8,1: it is in Period 3 (three shells) and Group I (one outer electron). Linking position to electron arrangement is the single most useful skill in this topic, because outer electrons decide bonding, charge of ions and reactivity.

Group I: the alkali metals

Group I contains lithium, sodium and potassium and is known as the alkali metals. They are soft enough to cut with a knife, have low densities (lithium, sodium and potassium float on water) and low melting points compared with most metals. Each atom has one outer electron, which it loses easily to form a singly charged positive ion such as Na+. They react vigorously with water to produce a metal hydroxide plus hydrogen gas, and the resulting solution is alkaline. Going down the group reactivity increases: lithium fizzes steadily, sodium melts into a ball and skates about, and potassium reacts so fast that the hydrogen ignites with a lilac flame. Reactivity rises down the group because the outer electron is further from the nucleus and more shielded, so it is lost more easily.

Group VII: the halogens

Group VII, the halogens, includes chlorine, bromine and iodine. They are coloured non-metals that exist as diatomic molecules such as Cl2. At room temperature chlorine is a pale green-yellow gas, bromine a red-brown liquid and iodine a grey-black solid, so colour darkens and physical state moves from gas to solid going down the group. Each atom needs to gain one electron to complete its outer shell, forming a singly charged negative ion such as Cl-. Unlike the alkali metals, halogen reactivity decreases down the group, because a larger atom attracts an incoming electron less strongly. This trend explains displacement reactions: a more reactive halogen will displace a less reactive one from a solution of its salt. For example, chlorine added to potassium bromide displaces bromine, turning the solution orange, because chlorine is above bromine in the group.

Group VIII: the noble gases

Group VIII (sometimes labelled Group 0) contains the noble gases such as helium, neon and argon. Their defining feature is that they are extremely unreactive, existing as single atoms rather than molecules. This inertness comes from their full outer electron shells: helium has two outer electrons and the others have eight, which is a very stable arrangement, so they have no tendency to gain, lose or share electrons. Their lack of reactivity makes them useful. Helium is much less dense than air and non-flammable, so it fills balloons and airships. Argon provides an inert atmosphere inside filament light bulbs and during welding to stop hot metal reacting with oxygen. Neon glows brightly when an electric current passes through it, so it is used in illuminated signs.

Transition elements

The block of metals in the centre of the table, between Group II and Group III, are the transition elements; common examples are iron, copper and zinc. Compared with Group I metals they are harder, stronger, denser and have much higher melting points, which makes them useful as construction materials and in machinery. They share several distinctive chemical properties: they often form ions with more than one charge (iron can form Fe2+ or Fe3+), their compounds are typically coloured (copper compounds are blue or green, many iron compounds are orange or brown), and both the metals and their compounds frequently act as catalysts. For instance, iron is the catalyst in the Haber process for making ammonia. These properties set them clearly apart from the highly reactive metals on the far left.

Predicting properties from position

Because the table is so ordered, you can estimate the properties of an element you have never met simply from its location. If an element sits in Group I, expect a soft, reactive metal that forms a 1+ ion; if it is at the bottom of Group VII, expect a dark solid non-metal that is less reactive than the halogens above it. Trends help you interpolate: melting point of a halogen, or reactivity of an alkali metal, can be sandwiched between its neighbours. Metallic character increases going down a group and decreases going across a period from left to right, so the most metallic elements are at the bottom left and the most non-metallic at the top right. Treating the table as a map of behaviour, rather than a list to memorise, is exactly what examiners want to see.