4.1.2.1,4.1.2.2,4.1.2.3 - The periodic table and its development

Groups, periods and electrons

The modern periodic table arranges elements by increasing atomic number and places elements with similar properties in vertical groups.

Atomic number is proton number. For neutral atoms it also gives electron number, so increasing atomic number builds a sequence of electronic structures. Similar outer-shell arrangements recur at regular intervals, giving the table its periodic pattern.

A group is a column and a period is a row. For the first twenty elements, the period number gives the number of occupied electron shells. For the main groups numbered 1 to 7 on the AQA table, the group number gives the number of outer-shell electrons. Group 0 is the full-outer-shell group: eight outer electrons except helium, whose first shell is full with two. Do not interpret Group 0 as meaning no electrons. The middle transition-metal block needs more detailed electron rules than these main-group shortcuts.

Magnesium, atomic number 12, has arrangement 2,8,2. Its three occupied shells put it in Period 3, and its two outer electrons put it in Group 2. Chlorine, atomic number 17, has 2,8,7, so it is also in Period 3 but in Group 7. Conversely, a Period 3, Group 1 element has three occupied shells and one outer electron; the first twenty-element filling rule gives 2,8,1, eleven electrons and therefore sodium.

Elements in a group have similar chemical properties because their atoms have the same number of outer electrons. Sodium and potassium both tend to lose their one outer electron. They are not identical substances: extra inner shells can change how easily that electron is lost, so reactivity can vary down a group.

Using position to predict chemical behaviour

The periodic table lets you make a reasoned prediction about an element before studying every one of its reactions.

Main-group metals towards the left have relatively few outer electrons and commonly lose them to form positive ions. Non-metals towards the right commonly gain electrons when reacting with metals, or share electrons when reacting with other non-metals. A full outer shell in a noble gas makes gaining, losing or sharing electrons difficult under ordinary conditions. These links explain chemical patterns instead of treating a group number as just a label.

For an illustrative prediction, suppose unfamiliar element X is a Group 1 metal below potassium. Since X has one outer electron, expect it to form a positive ion with charge +1 and react with water to form a hydroxide and hydrogen. Given the measured trend that Group 1 reactivity increases down the group, predict a more vigorous reaction than potassium under comparable conditions. The group relationship supports the reaction type; the trend supports the relative vigour. You cannot infer an exact reaction time from position alone.

A Group 7 non-metal has seven outer electrons and tends to gain one when reacting with a metal, forming a negative ion with charge −1. Given that reactivity decreases down this group, an element below bromine is expected to be less reactive than bromine. The opposite reactivity directions of Groups 1 and 7 come from losing versus gaining an electron, developed in the next lessons.

Predictions need an appropriate known pattern and comparable conditions. Boiling point is a physical property, while reactivity concerns chemical change: one trend does not automatically establish the other. An observed result that conflicts with a prediction is a reason to examine the conditions and model, not to ignore the result.

Mendeleev's table and tests of predictions

The modern arrangement grew from attempts to group known elements by their measured properties before subatomic particles were discovered.

Early scientists arranged elements by atomic weight, the historical term for their relative atomic mass. Their tables were incomplete because many elements had not been discovered. Strictly following increasing weight sometimes placed elements with different chemical properties in the same group.

Mendeleev left gaps where he expected undiscovered elements to belong. He used nearby elements to predict the missing elements' properties. He also changed the strict weight order in some places to preserve groups with similar properties. For example, iodine behaves as a halogen, so its chemical behaviour supported placing it with the halogens rather than in the preceding group merely to follow atomic weight order.

These decisions made testable predictions. When newly discovered elements filled the gaps and had the predicted properties, the observations supported his table. Successful predictions are stronger evidence than merely rearranging known facts to fit a diagram. An element with substantially different properties would have challenged a prediction or its underlying pattern. Communicating both predictions and measurements enabled other scientists to assess them.

Later knowledge of atomic structure explained the exceptions. Isotopes of one element have the same proton number but different masses, so average atomic masses need not increase strictly with proton number. Tellurium precedes iodine in the modern table because it has fewer protons, despite having a slightly greater average atomic mass. The modern table therefore uses atomic number, rather than weight. Mendeleev could not have used proton numbers to make his original decisions: protons had not yet been discovered.

Metals, non-metals and periodic position

Most elements are metals, found mainly on the left and towards the bottom of the periodic table, while non-metals are mainly towards the right and top.

Metals typically conduct electricity and heat well, have lustrous surfaces and can be shaped by bending or hammering rather than shattering. Most are solid at room temperature. Solid non-metals are often brittle and poor electrical conductors; several non-metals are gases. These are characteristic patterns with exceptions, not definitions that every element must satisfy. Graphite is a conducting non-metal form of carbon, and mercury is a liquid metal at room temperature.

The chemical distinction is that metals react to form positive ions by losing electrons. Non-metals do not form positive ions in the usual reactions considered here: they can form negative ions by gaining electrons or form covalent bonds by sharing. Sodium readily loses its one outer electron, whereas chlorine readily gains one to complete its outer shell. The transferred electron gives a positive sodium ion and a negative chloride ion; it does not turn one nucleus into a different element.

This relates position to structure. The left-hand main groups have few outer electrons, which can be lost to leave a full shell beneath. Towards the right, outer shells are closer to being full, so gaining or sharing electrons can achieve a stable arrangement. Group 0 already has full outer shells. Hydrogen is an exception to a simple left/right rule: despite its position above Group 1, it is a non-metal. Use the actual properties and electron behaviour alongside the table position.