Groups 1 & 7 and the Noble Gases

Groups, periods and the shape of the table

The Periodic Table arranges elements in order of increasing atomic number. The vertical columns are groups and the horizontal rows are periods.

The single most useful idea for this chapter is simple: the group number tells you the number of electrons in the outer shell. Group 1 elements have 1 outer electron, Group 7 have 7, and Group 0 (the noble gases) have a full outer shell. Because chemical reactions are all about electrons, elements in the same group behave alike.

Key terms Group — a vertical column; elements share the same number of outer-shell electrons and similar properties.

Period — a horizontal row; across it the elements change from metal to non-metal.

Outer shell — the highest occupied electron shell; it controls how an element reacts.

Group 1 — the alkali metals

Group 1 contains lithium (Li), sodium (Na) and potassium (K) — the alkali metals. They are unusual metals: soft enough to cut with a knife, with low densities (Li, Na and K all float on water) and low melting points compared with most metals. A freshly cut surface is shiny and silvery but quickly turns dull as it reacts with oxygen in the air.

Reaction with water. Alkali metals react vigorously with water to make a metal hydroxide plus hydrogen gas. The hydroxide is an alkali, which turns universal indicator blue/purple — this is where the name "alkali metal" comes from.

$$2Na + 2H_2O \rightarrow 2NaOH + H_2$$

Going down the group the reactions become more violent:

Lithium — floats and fizzes steadily, moving across the surface.
Sodium — melts into a ball from the heat released, whizzes about and fizzes rapidly.
Potassium — reacts so fast that the hydrogen produced ignites, burning with a lilac flame, and may spit and crack.

Worked example Write the equation for potassium reacting with water.

Potassium behaves like sodium: metal + water → metal hydroxide + hydrogen.

$$2K + 2H_2O \rightarrow 2KOH + H_2$$

Why does reactivity increase down the group? This is a guaranteed exam favourite, so learn the reasoning chain. To react, a Group 1 atom must lose its single outer electron.

1. Going down the group, atoms get larger — they have more electron shells.
2. So the outer electron is further from the nucleus.
3. There are also more inner shells, which shield the outer electron from the nuclear charge.
4. So the outer electron is less strongly attracted and is lost more easily.
5. Therefore the element is more reactive.

Exam tip Marks here are for the explanation, not just "K is more reactive than Na". A full answer must mention: more shells → outer electron further from nucleus → more shielding → electron lost more easily. Hit all three ideas.

Storage. Because they react with both water and oxygen, alkali metals are stored under oil. This keeps air and moisture away from the metal surface.

Group 7 — the halogens

Group 7 contains the halogens: chlorine (Cl), bromine (Br) and iodine (I). They are reactive non-metals that exist as diatomic molecules ($Cl_2$, $Br_2$, $I_2$). They are toxic and have sharp, choking smells.

Going down the group the halogens get darker in colour and their melting and boiling points increase (gas → liquid → solid).

Reactivity trend. Crucially, halogen reactivity decreases going down the group — the opposite direction to Group 1. To react, a halogen atom must gain one electron to complete its outer shell.

1. Going down the group, atoms get larger with more shells.
2. The outer shell is further from the nucleus and more shielded by inner electrons.
3. So an incoming electron is less strongly attracted.
4. The electron is gained less easily, so the halogen is less reactive.

Watch out Don't mix up the two trends. Group 1 loses an electron, so being big helps — reactivity goes up down the group. Group 7 gains an electron, so being big hinders — reactivity goes down down the group. So fluorine/chlorine are the most reactive halogens, while iodine is the least.

Displacement reactions of the halogens

A more reactive halogen displaces a less reactive halogen from a solution of its salt (a halide). This lets us rank reactivity experimentally.

Add chlorine water to potassium bromide solution. Chlorine is more reactive than bromine, so it pushes bromine out:

$$Cl_2 + 2KBr \rightarrow 2KCl + Br_2$$

The colourless solution turns orange-brown as bromine forms. Iodine being displaced gives a brown/purple colour. If you add iodine to potassium chloride, nothing happens — iodine is less reactive and cannot displace chlorine.

Worked example Chlorine is bubbled through potassium iodide solution and the mixture turns brown.

Chlorine is more reactive than iodine, so it displaces iodine. The brown colour is the iodine produced.

$$Cl_2 + 2KI \rightarrow 2KCl + I_2$$

Group 0 — the noble gases

Group 0 (also called Group 8) contains the noble gases: helium (He), neon (Ne), argon (Ar) and below them. They are colourless, monatomic gases that are almost completely unreactive (inert).

The reason is their electron arrangement: they have a full outer shell of electrons (helium has 2, the others have 8). A full outer shell is very stable, so noble gases have no tendency to lose, gain or share electrons — so they do not normally react.

Their inertness and other properties make them useful:

Helium — much less dense than air and non-flammable, so it is used in balloons and airships.
Argon — used to fill filament light bulbs; being unreactive, it stops the hot tungsten filament burning away. Also used as a shield gas in welding.
Neon — glows red-orange when a current passes through it, used in advertising signs.

Real world Hydrogen would lift a balloon even better than helium because it is lighter — but hydrogen is dangerously flammable (the Hindenburg airship disaster). Helium is chosen because it is safe and unreactive, a direct consequence of its full outer shell.

Periodic trends across a period

As you move across a period from left to right, elements change from metals to non-metals:

On the left (Groups 1–3) are reactive metals.
In the middle is a zig-zag dividing line.
On the right (Groups 5–7) are non-metals, ending in the unreactive Group 0.

So elements become less metallic and more non-metallic across a period, and more metallic going down a group. This is why the most reactive metals sit at the bottom-left (e.g. potassium) and the most reactive non-metals near the top-right (e.g. fluorine/chlorine).

Exam tip Two trends to memorise and never confuse:

Going down Group 1 → metals get more reactive.

Going down Group 7 → non-metals get less reactive.

Both are explained by the same idea — bigger atoms, more shielding, outer electron further from the nucleus — but it helps a metal lose an electron and hinders a non-metal gaining one.