Electronic Configuration & Reactivity

An element's chemical reactivity is decided almost entirely by its outer-shell electrons — how many there are, and how easily they can be lost or gained.

• Metals (like Group 1) react by LOSING outer-shell electron(s) to form positive ions.
• Non-metals (like Group 7) react by GAINING electron(s) to form negative ions.
• The atom wants to reach a full (stable) outer shell — the arrangement that the noble gases (Group 0) already have.

So "more reactive" means:

• for a metal, the outer electron is lost more easily;
• for a non-metal, an electron is gained more easily.

Two atomic-structure ideas explain every trend in this topic:

1. Distance from the nucleus. Going down any group, each element has one more electron shell, so the atom is larger and the outer shell is further from the positive nucleus. The further away an electron is, the weaker the electrostatic attraction between it and the nucleus.
2. Shielding (screening). The inner-shell electrons sit between the nucleus and the outer shell. They repel the outer electrons and block some of the nuclear pull. More inner shells = more shielding = weaker net attraction on the outer shell.

The trick: whichever group you are in, ask "does a weaker pull on the outer shell make it EASIER or HARDER to react?" — that single question gives you the trend.

Group 1 elements (alkali metals): Li, Na, K, Rb, Cs. Each has just 1 outer-shell electron and reacts by LOSING it to form a +1 ion:

M → M⁺ + e⁻

Reactivity trend: INCREASES down the group. Li (least reactive) < Na < K < Rb < Cs (most reactive). You can see this with water — Li fizzes gently, Na melts into a ball that zooms about, K ignites with a lilac flame, and Cs reacts explosively.

Explanation that earns the marks. Going down Group 1:

1. Each element has one more electron shell (Li has 2 shells, Na 3, K 4…), so the atom gets larger.
2. The outer electron is further from the nucleus → the attraction between the nucleus and the outer electron is weaker.
3. There is more shielding from the extra inner shells, which also reduces the pull on the outer electron.
4. So the outer electron is lost more easily → the metal is MORE reactive.

Watch out (extension). Going down the group the nucleus also gains more protons, which on its own would pull harder on the outer electron. But the effects of greater distance and more shielding outweigh this, so the net attraction gets weaker and reactivity rises.

Other Group 1 facts to bank:

• Soft (cut with a knife), low density (Li, Na, K float on water), low melting points.
• React with water to give a metal hydroxide + hydrogen: 2Na + 2H₂O → 2NaOH + H₂.
• Their hydroxide solutions are alkaline — hence the name "alkali metals".

Group 7 elements (halogens): F, Cl, Br, I. Each has 7 outer-shell electrons and reacts by GAINING 1 electron to form a −1 ion:

X + e⁻ → X⁻

The halogens also exist as diatomic molecules (F₂, Cl₂, Br₂, I₂). Useful states and colours:

Halogen	State at room temp	Colour
Cl₂	Gas	Pale green
Br₂	Liquid	Orange / red-brown
I₂	Solid	Dark grey (purple vapour)

Reactivity trend: DECREASES down the group. F (most reactive) > Cl > Br > I.

Explanation that earns the marks. Going down Group 7:

1. Each halogen has one more electron shell (F has 2 shells, Cl 3, Br 4…), so the atom gets larger.
2. The outer shell — where the incoming electron must go — is further from the nucleus → the attraction on an incoming electron is weaker.
3. There is more shielding from the extra inner shells, further reducing the pull on an incoming electron.
4. So an electron is gained less easily → the halogen is LESS reactive.

(Notice it is the same two ideas as Group 1 — distance and shielding — but because halogens gain rather than lose, a weaker pull makes them less reactive, not more.)

Halogen displacement reactions. A more reactive halogen will DISPLACE a less reactive halide from a solution of its salt:

Cl₂ + 2KBr(aq) → 2KCl(aq) + Br₂(aq) (colourless → orange-brown)

Cl₂ + 2KI(aq) → 2KCl(aq) + I₂(aq) (colourless → dark brown)

Br₂ + 2KI(aq) → 2KBr(aq) + I₂(aq)

The reverse never happens — a less reactive halogen cannot displace a more reactive halide (e.g. I₂ + KBr → no reaction).

Displacement is therefore a neat test of reactivity order: if a reaction happens (a colour change appears), the halogen you added is more reactive than the halide it displaced.

Group 0 elements (noble gases): He, Ne, Ar, Kr, Xe, Rn. They exist as single, unbonded (monatomic) atoms — not diatomic molecules.

Why they are unreactive — the one idea you must state. Noble gases have a FULL OUTER ELECTRON SHELL, which is a very stable arrangement:

Element	Electron configuration	Outer-shell electrons
He	2	2 (full 1st shell)
Ne	2, 8	8 (full)
Ar	2, 8, 8	8 (full)

Reactions happen because atoms are trying to reach a full outer shell (Group 1 metals lose an electron, halogens gain one). Noble gases already have a full outer shell, so:

• they do not need to lose electrons,
• they do not need to gain electrons,
• they do not need to share electrons.

With no driving force to react, they stay as separate atoms and do not form compounds under normal conditions → they are unreactive (inert).

Uses (which all rely on being unreactive):

• Helium — balloons and airships (low density, and unlike hydrogen it will not burn).
• Argon — fills filament light bulbs (stops the hot filament reacting/oxidising); shielding gas in welding.
• Neon — glowing advertising signs.