The periodic table and group trends (C4.1)

The periodic table provides a framework for predicting how elements behave. OCR Gateway A typically asks students to predict reactions based on position in the periodic table and the reactivity series.

Structure of the periodic table

• Elements are arranged in order of increasing atomic number.
• Periods (horizontal rows): elements with the same number of electron shells.
• Groups (vertical columns): elements with the same number of outer-shell electrons → similar chemical properties.
• Metals are on the left and centre; non-metals on the right.

Group 1 — the alkali metals (Li, Na, K, Rb, Cs, Fr)

All have 1 outer electron; they react by losing it to form M⁺ ions.

Reactions with water:

metal + water → metal hydroxide + hydrogen
2Na + 2H₂O → 2NaOH + H₂

Trend down the group: reactivity INCREASES.

• Because the outer electron is further from the nucleus / greater shielding by inner shells → easier to lose → more reactive.

Observations: lithium fizzes gently; sodium fizzes vigorously; potassium ignites and burns with a lilac flame.

All group 1 hydroxides form alkaline solutions (the clue is in the name "hydroxide").

Group 7 — the halogens (F, Cl, Br, I, At)

All have 7 outer electrons; they react by gaining one electron to form X⁻ ions (halide ions).

Trend down the group: reactivity DECREASES.

• Because added electron is harder to gain when further from the nucleus / less attracted / more shielding.

Displacement reactions

A more reactive halogen displaces a less reactive halide from solution:

chlorine + sodium bromide → sodium chloride + bromine
Cl₂ + 2NaBr → 2NaCl + Br₂

Chlorine is more reactive than bromine → it takes the electron away from Br⁻.

The solution turns from colourless to orange-brown (bromine in solution).

Physical states (at room temperature):

• Fluorine (F₂) — gas (pale yellow-green)
• Chlorine (Cl₂) — gas (yellow-green)
• Bromine (Br₂) — liquid (red-brown)
• Iodine (I₂) — solid (grey-purple; violet vapour)

Melting/boiling points increase down the group.

Group 0 — the noble gases (He, Ne, Ar, Kr, Xe, Rn)

• Full outer shell → stable configuration → extremely unreactive.
• Exist as monatomic gases.
• Boiling points and densities increase down the group.

Uses: helium (balloons, breathing gas for deep-sea divers), argon (inert atmosphere in welding/filament bulbs), neon (lighting signs).

Transition metals

The central block of the periodic table. Key properties:

• Good conductors of electricity and heat.
• Higher melting points than Group 1 metals.
• Denser and stronger than Group 1 metals.
• Can form ions with multiple charges (e.g. Fe²⁺, Fe³⁺; Cu²⁺, Cu⁺).
• Often form coloured compounds (copper sulfate is blue; iron(III) chloride is orange-brown).
• Many are catalysts (iron in the Haber process; vanadium oxide in the Contact process).

Reactivity series (metals)

From most to least reactive:

K > Na > Ca > Mg > Al > Zn > Fe > Ni > Sn > Pb > H > Cu > Ag > Au > Pt

(Mnemonic: "Kicky Naughty Cows Make A Zillion Frenzied Nations Slay Lions, Houses, Cats, And Go Potty")

A more reactive metal displaces a less reactive metal from its salt solution:

zinc + copper sulfate → zinc sulfate + copper
Zn + CuSO₄ → ZnSO₄ + Cu

Metals above hydrogen in the series react with dilute acid to produce hydrogen gas; those below do not.

Common Gateway-paper mistakes

1. Saying reactivity increases down Group 7 — it DECREASES (the opposite of Group 1).
2. Forgetting to explain displacement reactions in terms of relative reactivity.
3. Saying transition metals only form one type of ion.
4. Confusing periods (horizontal) with groups (vertical).
5. Forgetting that noble gases (Group 0) are monatomic — they don't form diatomic molecules.