Periodic Table Trends

Summary: Elements are arranged by increasing atomic number. Periods are horizontal rows; groups are vertical columns. Metallic character decreases across a period and increases down a group. Tags: igcse chemistry periodic-table Created: 2026-07-14 Last Updated: 2026-07-16

Arrangement of the Periodic Table

The modern Periodic Table arranges elements in order of increasing atomic number (number of protons).

• Groups are vertical columns. Elements in the same group have the same number of electrons in their outer shell, which gives them similar chemical properties.
• Periods are horizontal rows. Elements in the same period have the same number of occupied electron shells.

The Periodic Table can be used to predict the properties of elements based on their position.

Group Number and Outer-Shell Electrons

For Groups 1-2 and 13-18 (the main groups), the group number tells the number of electrons in the outer shell:

Period number = number of occupied electron shells:

• Period 1: 1 shell (H, He)
• Period 2: 2 shells (Li to Ne)
• Period 3: 3 shells (Na to Ar)
• Period 4: 4 shells (K to Kr)

Why elements in the same group have similar chemical properties: Chemical reactions involve the gain, loss, or sharing of outer-shell electrons. Since elements in the same group have the same number of outer-shell electrons, they react in similar ways. For example, all Group 1 metals form 1+ ions; all Group 7 halogens form 1- ions.

Metals vs Non-Metals

Position on the Periodic Table:

• Metals are found on the left and centre of the Periodic Table (Groups 1-2, the transition metal block, and to the left of the “staircase” line starting at boron). Approximately 75% of elements are metals.
• Non-metals are found on the right of the Periodic Table (upper right, above and to the right of the staircase). The staircase line runs from boron (B) down through silicon (Si), arsenic (As), tellurium (Te), and astatine (At).

Trend in metallic character:

• Down a group: metallic character increases (the elements become more metallic). Outer electrons are further from the nucleus and more easily lost.
• Across a period (left to right): metallic character decreases (elements change from metallic to non-metallic).

Typical properties of metals:

• High melting and boiling points (most, but not all — e.g., mercury is liquid at RTP, and Group 1 metals have relatively low melting points)
• Good conductors of heat and electricity
• Malleable (can be hammered into sheets) and ductile (can be drawn into wires)
• Shiny (lustrous) when freshly cut or polished
• Form basic oxides (e.g., Na2O, MgO, CaO) which react with acids to form salt + water
• Lose electrons to form positive ions (cations) in reactions

Typical properties of non-metals:

• Generally lower melting and boiling points (many are gases at RTP)
• Poor conductors of heat and electricity (except graphite, a form of carbon)
• Brittle when solid (not malleable or ductile)
• Dull appearance (not shiny)
• Form acidic oxides (e.g., SO2, CO2, P4O10) which react with alkalis to form salt + water (or neutral oxides like CO and H2O)
• Gain or share electrons in reactions

Silicon (Si) is found along the staircase line and is a metalloid (semi-metal). It has some properties of metals (shiny, semi-conductor of electricity) and some of non-metals (brittle, forms an acidic oxide SiO2).

Trends Across Period 3 (Na to Ar)

Period 3 is the best example for illustrating periodic trends. It runs from sodium (Na, Z=11) to argon (Ar, Z=18).

Melting point trend across Period 3:

• Na, Mg, Al: metallic bonding — strong electrostatic attraction between positive metal ions and delocalised electrons. Melting points increase from Na to Al because the metal ions have a greater charge (Na+, Mg2+, Al3+) and more delocalised electrons per atom, giving stronger metallic bonds.
• Si: giant covalent structure — very high melting point (1410 °C) because strong covalent bonds must be broken throughout the entire lattice.
• P4, S8, Cl2, Ar: simple molecular structures — low melting points because only weak intermolecular forces need to be overcome (not the strong covalent bonds within molecules). Ar is monatomic with the weakest forces of all.

Electrical conductivity across Period 3:

• Na, Mg, Al: good conductors (delocalised electrons in metallic bonding can move freely and carry charge).
• Si: semi-conductor (some conductivity, but electrons are held in covalent bonds so less free than in metals).
• P, S, Cl, Ar: non-conductors (no delocalised electrons or free ions).

Nature of oxides across Period 3:

• Na2O, MgO: basic oxides — react with acids to form salt + water. Na2O + 2HCl → 2NaCl + H2O. MgO + 2HCl → MgCl2 + H2O.
• Al2O3: amphoteric oxide — reacts with BOTH acids AND alkalis. Al2O3 + 6HCl → 2AlCl3 + 3H2O. Al2O3 + 2NaOH + 3H2O → 2NaAl(OH)4.
• SiO2: acidic oxide (but insoluble in water; reacts with hot concentrated alkali). SiO2 + 2NaOH → Na2SiO3 + H2O.
• P4O10, SO2/SO3: acidic oxides — react with water to form acidic solutions. P4O10 + 6H2O → 4H3PO4. SO2 + H2O → H2SO3.
• Cl2O7: acidic oxide. Ar: no oxide (unreactive).

Example: Explain why the melting point of aluminium (660 °C) is much higher than that of sodium (98 °C).

• Both Na and Al have metallic bonding.
• Al forms Al3+ ions; Na forms Na+ ions.
• Al has three delocalised electrons per atom; Na has only one.
• The electrostatic attraction between Al3+ and the delocalised electrons is stronger than between Na+ and delocalised electrons, so more energy is needed to overcome the metallic bonds in Al.

Trends Down Groups

Atomic radius:

• Atomic radius increases down any group because each element has one more occupied electron shell than the one above it. The outer electrons are further from the nucleus.

Reactivity in Groups 1 and 2 (metals):

• Reactivity increases down the group.
• Explanation: As you go down the group, the outer electron is in a shell further from the nucleus and there is more shielding by inner electron shells. The attraction between the nucleus and the outer electron is weaker, so the outer electron is more easily lost.

Reactivity in Group 7 (non-metals):

• Reactivity decreases down the group.
• Explanation: As you go down the group, the outer shell is further from the nucleus and there is more shielding. The attraction for an incoming electron is weaker, so it is harder to gain an electron.

Melting points down groups:

• For Group 1 metals: melting points generally decrease down the group (Li: 181 °C, Na: 98 °C, K: 63 °C). The metallic bonds get weaker because the ions get larger, reducing the electrostatic attraction to delocalised electrons.
• For Group 7 halogens: melting and boiling points increase down the group (F2 and Cl2 are gases, Br2 is liquid, I2 is solid at RTP). The molecules get larger with more electrons, so intermolecular forces get stronger.

The Diagonal Relationship

Elements in Group 1 and Group 2, when compared with transition metals one period down, show a diagonal relationship where properties become more similar. For example:

• Li (Group 1, Period 2) has some similarities to Mg (Group 2, Period 3): both form nitrides when heated in nitrogen, both have relatively high melting points compared to other Group 1 metals, both carbonates decompose easily on heating.
• Be (Group 2, Period 2) and Al (Group 13, Period 3) both form amphoteric oxides.

This reflects counterbalancing trends: going right across a period decreases metallic character, while going down a group increases it. Moving diagonally (down-left) keeps metallic character roughly similar.

Key Points

• The Periodic Table is arranged by increasing atomic number (not atomic mass)
• Group = vertical column; elements in the same group have the same number of outer-shell electrons
• Period = horizontal row; elements in the same period have the same number of occupied electron shells
• Group number = number of outer-shell electrons (Groups 1-2 and 13-18)
• Period number = number of occupied shells
• Metals: left side; Non-metals: right side
• Metallic character decreases across a period; increases down a group
• Metals form basic oxides; non-metals form acidic oxides; Al2O3 is amphoteric
• Atomic radius increases down a group (more electron shells)
• Group 1: reactivity increases down; Group 7: reactivity decreases down

Key Concepts from Past Papers

• Same number of outer-shell electrons explains similar chemical properties in a group
• Atomic radius increases down a group due to additional electron shells

Keywords from Past Papers

group, protons, electrons, hydrogen, carbon, cooh, down, metal, circle, around, decreases, mixture, neutrons, properties, elements

Related Notes

• Group 1 Alkali Metals
• Group 7 Halogens
• Group 0 Noble Gases
• Transition Metals
• Atomic Structure

Sources

• OpenStax Chemistry 2e — Chapter 6: Electronic Structure and Periodic Properties, Rice University (free, CC BY 4.0)
• BBC Bitesize GCSE Chemistry — Periodic Table Trends, BBC (free educational resource)
• Cambridge IGCSE Chemistry 0620 — Syllabus topic 8 (The Periodic Table), Cambridge Assessment International Education
• CK-12 Chemistry for High School — Periodic Trends, CK-12 Foundation (free, CC BY-NC 3.0)

Past Paper Sources

• 0620/31 May/June 2015: Q22(b)(iii) (0m), Q33(b)(ii) (0m), Q44(a)(i) (0m) (+2 more)
• 0620/32 Feb/March 2015: Q11(b)(ii) (3m), Q66(b)(i) (1m)
• 0620/32 Feb/March 2017: Q33(b)(i) (1m), Q77(c)(i) (1m)
• 0620/32 Feb/March 2018: Q22(b)(i) (1m), Q33(b)(ii) (1m), Q55(d)(i) (1m)
• 0620/32 Feb/March 2019: Q22(a)(i) (2m)
• 0620/32 Feb/March 2020: Q44(a)(i) (1m)
• 0620/32 Feb/March 2021: Q33(b)(ii) (1m), Q66(a)(i) (1m)
• 0620/32 Feb/March 2022: Q33(b)(i) (1m)
• 0620/32 Feb/March 2023: Q66(a)(i) (1m)
• 0620/32 May/June 2018: Q22(c)(ii) (2m), Q33(a)(ii) (1m), Q77(a)(iii) (1m) (+1 more)
• 0620/32 May/June 2020: Q44(a)(i) (1m)
• 0620/33 May/June 2016: Q11(b)(i) (1m), Q33(a)(ii) (1m)

Common Misconceptions