The Periodic Table: Structure and Trends

The Periodic Table arranges elements by proton number, so that elements with similar properties fall in the same column. Once you can read a position, you can predict the chemistry.

---

1. How it is arranged

Elements are arranged in order of increasing PROTON NUMBER.

GROUPS are the vertical COLUMNS. All elements in a group have the SAME number of OUTER-SHELL ELECTRONS, which is why they react similarly. PERIODS are the horizontal ROWS. The period number equals the number of occupied electron shells.

Groups go down, periods go across. Confusing group numbers with period numbers was recorded, as was general confusion about the arrangement.

Reading position from electronic configuration — a standard question:

Group number = number of OUTER-shell electrons. Period number = number of SHELLS.

Example: an element with configuration 2,8,6 is in Group VI, Period 3.

Tutors flagged that questions often ask for group and period from the electronic configuration — practise both directions.

Predict the valence electrons from the group number — also flagged directly.

---

2. The named groups

Group	Name	Outer electrons
I	alkali metals	1
II	alkaline earth metals	2
VII	HALOGENS	7
VIII (0)	NOBLE GASES	8 (helium has 2)
middle block	transition elements	variable

Group VII are the HALOGENS; Group VIII are the NOBLE GASES. Both names were recorded as forgotten or garbled — “Madagens” for halogens was one attempt. They are worth learning as vocabulary.

Transition elements sit in the middle block, between Groups II and III. Forgetting where Group I and the transition metals sit was recorded.

---

3. Metals and non-metals

METALS are on the LEFT and in the middle. NON-METALS are on the RIGHT. A staircase line separates them.

	Metals	Non-metals
Conduct heat and electricity	yes	no (except graphite)
Appearance	shiny, malleable, ductile	dull, brittle when solid
Melting point	usually high	usually low
Ions formed	positive (lose electrons)	negative (gain electrons)
Oxides	basic	acidic

Two liquid elements at room temperature: MERCURY (a metal) and BROMINE (a non-metal). Both were recorded as forgotten — bromine was even transcribed as “Berlin”. They come up in multiple-choice questions regularly.

Metalloids (e.g. silicon) sit along the staircase and show intermediate properties.

---

4. Trends ACROSS a period

Going left to right across a period:

The number of protons INCREASES, so the NUCLEAR CHARGE increases. Electrons are added to the SAME shell, so the shielding stays roughly constant. The nucleus attracts the outer electrons MORE STRONGLY, so the ATOMIC RADIUS DECREASES.

Give BOTH points when explaining atomic size across a period: the increasing nuclear charge AND that electrons are added to the same principal shell. Tutors flagged this as a two-part mark.

Only PROTONS contribute to nuclear charge — electrons do not. A recorded misunderstanding had electrons contributing. Adding electrons to the same shell does not shield the outer electrons much, which is why the atom gets smaller, not larger.

Also across a period: elements change from metals to non-metals, and oxides change from basic to acidic.

---

5. Trends DOWN a group

Going down a group:

More electron shells are added, so the ATOMIC RADIUS INCREASES. The outer electrons are FURTHER from the nucleus and better SHIELDED by inner shells, so the force of attraction is WEAKER.

The consequences depend on whether the group loses or gains electrons:

Group I (metals, LOSE an electron): reactivity INCREASES down the group, because the outer electron is more easily lost. Group VII (non-metals, GAIN an electron): reactivity DECREASES down the group, because it is harder to attract an extra electron.

Reactivity DECREASES down Group VII. A recorded error had it increasing — the two groups trend in opposite directions, and this is the single most examined trend on the page.

Weaker attraction explains BOTH trends. A student was recorded as confused about why less force of attraction leads to more violent reactions — for a metal, weaker attraction means the electron is lost more easily, so the reaction is more vigorous. For a halogen, weaker attraction means it captures an electron less easily, so it is less reactive.

---

6. The noble gases (Group VIII)

The noble gases are UNREACTIVE (inert) because they have a FULL OUTER SHELL — 8 electrons (or 2 for helium) — so they have no tendency to gain, lose or share electrons.

They are monatomic (existing as single atoms) and colourless.

Uses: helium in balloons (low density, non-flammable), argon in light bulbs and welding (inert atmosphere), neon in lighting.

A full outer shell is the reason for their inertness — and it is also the target every other element is trying to reach when it bonds.

---

7. Explaining melting and boiling point trends

Use precise keywords. Tutors flagged this — melting point explanations need the structure and the forces.

• Group I: melting point decreases down the group — the metallic bonding weakens as atoms get larger
• Group VII: melting point increases down the group — larger molecules have stronger intermolecular forces
• Across a period: rises to a peak at the giant covalent element (carbon or silicon), then falls sharply for the simple molecular non-metals

---

8. Mistakes that cost marks

Confusing groups with periods.

Forgetting the group names — halogens, noble gases.

Saying reactivity increases down Group VII.

Giving only one reason for the atomic size trend.

Saying electrons contribute to nuclear charge.

Forgetting mercury and bromine as the liquid elements.

Saying atomic size increases across a period.

Explaining noble gas inertness without “full outer shell”.

Misreading group or period from a configuration.

---

Frequently asked questions

How is the Periodic Table arranged? By increasing proton number.

What do elements in the same group share? The same number of outer-shell electrons.

What does the period number tell you? The number of electron shells.

Which group are the halogens? Group VII. The noble gases are Group VIII.

Which elements are liquid at room temperature? Mercury (metal) and bromine (non-metal).

Why does atomic radius decrease across a period? Increasing nuclear charge with electrons added to the same shell.

Why does reactivity increase down Group I? The outer electron is further away and better shielded, so it is lost more easily.

Why does reactivity decrease down Group VII? It is harder to attract an extra electron into a larger, more shielded atom.

Why are noble gases unreactive? They have a full outer shell.

How do I find group and period from a configuration? Outer electrons = group; number of shells = period.

---

Quick revision checklist

• I know the table is ordered by proton number
• I know groups are columns and periods are rows
• I can find group and period from an electronic configuration
• I know the names of Groups I, VII and VIII
• I know where the transition elements sit
• I can distinguish metals from non-metals and their properties
• I know mercury and bromine are the liquid elements
• I can explain the atomic size trend across a period with both reasons
• I know only protons affect nuclear charge
• I can explain the size trend down a group
• I know reactivity increases down Group I and decreases down Group VII
• I can explain why each trend goes that way
• I can explain noble gas inertness and give their uses

---

These notes cover the structure and trends of the Periodic Table in the Cambridge IGCSE Chemistry (0620) syllabus and are written for Grade 9–11 / Year 10–11 students. They are based on teaching patterns observed across a large set of one-to-one IGCSE Chemistry lessons, with particular attention to the errors students make most often and the wording examiners reward. Always check the current syllabus and data booklet for your own exam series.