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The periodic table is the single most important organising idea in chemistry: a chart that arranges all the elements so that those with similar properties fall into the same column. Its real power is that the elements in a group behave in related, predictable ways, and the changes in their behaviour as you go down a group follow clear trends — every one of which can be explained by electronic structure. This lesson, part of Topic C2 of your OCR Gateway Combined Science course, first traces how the table was built (Newlands and Mendeleev) and why it is now ordered by atomic number, then examines the three groups you must know in detail: Group 1 (the alkali metals), Group 7 (the halogens) and Group 0 (the noble gases).
By the end of this lesson you should be able to describe how the periodic table developed, explain why Mendeleev left gaps and why the modern table is ordered by atomic number, use an element's position to deduce its properties, describe the reactions and trends of the alkali metals and halogens, and explain why the noble gases are unreactive.
This lesson develops AO1 (recalling the development of the table and the reactions of the alkali metals and halogens) and AO2 (using an element's position to deduce its properties), reaching AO3 when you explain the Group 1 and Group 7 reactivity trends in terms of electronic structure.
In the early 1800s the structure of the atom was unknown, so the only property chemists could measure to compare elements was their relative atomic weight. The first attempts to organise the elements simply put them in order of increasing atomic weight.
In 1864 John Newlands noticed that, when the known elements were listed in order of atomic weight, every eighth element seemed to have similar properties — rather like the repeating notes of a musical scale. He called this the law of octaves. It was a genuine insight that properties repeat periodically, and it worked for the lighter elements. But it had a serious flaw: Newlands left no gaps for elements not yet discovered, so he was forced to cram known elements into fixed slots, sometimes placing very different elements in the same group. The pattern also broke down for the heavier elements, so his ideas were not widely accepted.
In 1869 Dmitri Mendeleev produced a table that overcame these problems and is recognised as the ancestor of the modern one. He, too, began by ordering the elements by atomic weight, but he did two things that made his table work where Newlands' had failed.
The most powerful step was using the gaps to predict the properties of undiscovered elements. For the gap below silicon he described a missing element he called "eka-silicon", predicting its atomic weight, its density and the formula of its oxide. When germanium was discovered in 1886, its properties matched his predictions astonishingly well. These successful predictions were exactly what convinced chemists his table was correct — a table that predicts is far more than a list.
Exam Tip: Mendeleev's genius was not just listing elements by weight — it was leaving gaps for undiscovered elements and predicting their properties. When those elements (such as germanium) were found and matched, the predictions were strong evidence for his table.
Ordering by atomic weight left a few anomalies — pairs of elements that seemed to be in the "wrong" order. The explanation came in the early twentieth century, once the structure of the atom was understood. The discovery of protons showed that each element has a fixed number of protons — its atomic number — and the discovery of isotopes (atoms of the same element with different numbers of neutrons, and so different masses) explained why atomic weight did not always increase smoothly. When the elements are ordered by atomic number instead of atomic weight, the anomalous pairs fall into their correct places automatically, with no swapping needed.
A classic example is argon and potassium: by atomic weight, argon (about 40) is heavier than potassium (about 39), which would reverse them — but by atomic number argon (18) correctly comes before potassium (19), placing argon with the noble gases and potassium with the reactive metals, matching their properties.
Exam Tip: The modern table is ordered by atomic (proton) number, not atomic weight. Ordering by atomic number resolves the anomalies (like argon/potassium) that forced Mendeleev to swap pairs, because isotopes make atomic weight an imperfect guide. A common misconception is to say the modern table is "ordered by mass".
The modern table arranges the elements into groups and periods, and this layout links directly to electronic structure (which you met in C1).
The table also divides into metals and non-metals. Metals are found on the left and centre (the large majority of elements); non-metals are on the right-hand side, separated by a "staircase" line with the metals to its left.
An element X is in Group 1, Period 3. Use its position to state how many outer-shell electrons it has, how many shells, whether it is a metal or non-metal, and one chemical property.
Step 1 — outer electrons = group number = 1.
Step 2 — number of shells = period number = 3.
Step 3 — Group 1 is on the left, so X is a metal.
Step 4 — Group 1 metals (the alkali metals) are very reactive and react with water; with 1 outer electron, X forms a +1 ion.
Answer: X has 1 outer electron, 3 shells, is a reactive metal (an alkali metal), and forms a +1 ion. (This element is sodium.)
Exam Tip: The group number = outer-shell electrons and the period number = number of shells. From those two facts you can read off an element's likely charge and whether it is a metal — a quick, reliable source of marks.
Group 1 elements — lithium, sodium, potassium and below — are the alkali metals. They each have one electron in the outer shell, and they share a distinctive set of properties:
Reaction with water. The alkali metals react vigorously with water to produce a metal hydroxide (an alkali — hence the group's name) and hydrogen gas:
sodium+water→sodium hydroxide+hydrogen 2Na+2H2O→2NaOH+H2
The metal floats, fizzes and skates about on the surface as hydrogen is released. Lithium reacts steadily; sodium reacts faster and may melt into a ball; potassium reacts so vigorously that the hydrogen ignites with a lilac flame. This shows the trend clearly: reactivity increases down Group 1.
Explaining the trend. Going down the group, each element has one more electron shell, so the outer electron is further from the nucleus and is shielded from the nuclear charge by more inner shells. Both effects mean the outer electron is less strongly attracted to the nucleus, so it is lost more easily. Since reacting involves losing that outer electron, the metals get more reactive down the group.
Exam Tip: Reactivity increases down Group 1. The reason to write is: down the group the outer electron is further from the nucleus and more shielded, so it is lost more easily — and losing the outer electron is what reacting requires.
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