Difficulty: Intermediate | Prerequisites: Atomic Structure and Quantum Mechanics notes, basic familiarity with the periodic table layout.
The periodic table is the single most useful organising tool in chemistry. Every element's position on it is determined by its electron configuration, and that position predicts the element's physical and chemical properties. This topic covers how the table is structured (groups, periods, blocks), the key families of elements, and the six periodic trends you need to know: atomic radius, ionic radius, ionization energy, electronegativity, electron affinity and reactivity. If you understand electron configurations from the Atomic Structure notes, you are ready.
Valence electrons determine an element's group and its chemistry. Across a period (left to right), atomic radius shrinks, ionization energy and electronegativity climb, and metals give way to nonmetals. Down a group, the opposite happens because each row adds another electron shell. Metals are most reactive at the bottom-left of the table; nonmetals are most reactive at the top-right.
Valence electrons
The electrons in an atom's outermost shell. They determine the element's position on the periodic table and control how it bonds.
Period
A horizontal row on the periodic table. All elements in the same period share the same principal energy level (same number of electron shells).
Group (family)
A vertical column on the periodic table. Elements in the same group have the same number of valence electrons and similar chemical behaviour.
s-block elements
Groups 1 and 2 (plus helium). Their valence electrons occupy s orbitals. Includes the alkali metals (Group 1) and alkaline earth metals (Group 2).
p-block elements
Groups 13 to 18. Their valence electrons fill p orbitals. The p-block includes metals, metalloids, nonmetals and noble gases.
d-block elements (transition metals)
Groups 3 to 12. Their distinguishing electrons fill d orbitals. More similar to one another across a period than s- or p-block elements are.
f-block elements (inner transition metals)
The lanthanides and actinides. Their distinguishing electrons fill f orbitals. Lanthanides are silvery with high melting points; actinides are radioactive.
Atomic radius
Half the distance between the nuclei of two bonded atoms of the same element. In simple terms, how big the atom is.
Ionic radius
The radius of an ion (a charged atom). Cations are smaller than their parent atoms; anions are larger.
Ionization energy (IE)
The energy required to remove the outermost electron from a gaseous atom. Think of it as how tightly the atom holds onto its electrons.
Electronegativity
A measure of how strongly an atom attracts electrons in a chemical bond. Fluorine is the most electronegative element.
Electron affinity
The energy change when a gaseous atom gains an electron. In simple terms, how much an atom "wants" an extra electron.
Alloy
A mixture of metals (or a metal with a nonmetal) whose composition can be adjusted to produce different properties. Can be interstitial (small atoms fit in gaps) or substitutional (atoms swap places).
Diagonal relationship
Similarities in properties between a period-2 element and the period-3 element one group to its right (e.g. Li and Mg, Be and Al).
Valence electrons determine an element's group (column) and its chemistry.
Elements in the same group share similar properties, but they are not identical because they differ in their number of inner electrons.
Elements in the same period share the same principal energy level (number of shells).
Diagonal relationships exist between period-2 and period-3 elements sitting diagonally (e.g. Li/Mg, Be/Al). These pairs show unexpectedly similar properties.
Hydrogen: unique. Can act as both a metal and a nonmetal (gains or loses electrons).
Alkali metals (Group 1): most reactive metals. Sodium and potassium are the most abundant. Reactivity increases going down the group.
Alkaline earth metals (Group 2): very reactive. Calcium is essential for bones and teeth (found in calcium carbonate, which decomposes to lime). Magnesium is used in lightweight alloys and is essential for metabolism, photosynthesis and muscle function.
Aluminium: the most abundant metal in Earth's crust.
Carbon: forms millions of organic compounds. Graphite and diamond are allotropes (same element, different structures).
Silicon and oxygen: the most abundant elements in Earth's crust, found in silica (which melts and cools to form glass).
Nitrogen: combines with hydrogen to form ammonia (cleaning products). Nitric acid (from ammonia) is used in fertilisers, explosives and dyes.
Phosphorus: phosphates in fertilisers and cleaning products harm the environment.
Sulfur: sulfur dioxide reacts with water to form one of the acids in acid rain.
Halogens (Group 17): extremely reactive nonmetals. Many plastics contain chlorine. Silver bromide and iodide coat photographic film.
Noble gases (Group 18): stable and largely unreactive. Used in blimps, neon lights and diving tanks.
Transition metals (d-block): more similar across a period than s- or p-block elements. The more unpaired d electrons, the harder the metal and the higher its melting and boiling points. Ferromagnetic metals have ions permanently aligned in a magnetic field. Many are strategic materials.
Lanthanides: silvery, high melting points, found mixed in nature, hard to separate.
Actinides: radioactive.
Property | Across period (left to right) | Down group (top to bottom) |
|---|---|---|
Atomic radius | Decreases | Increases |
Ionic radius | Decreases | Increases |
Ionization energy | Increases | Decreases |
Electronegativity | Increases | Decreases |
Electron affinity | Increases | Decreases |
Metal reactivity | Decreases | Increases |
Nonmetal reactivity | Increases | Decreases |
Across a period: decreases left to right. More protons in the nucleus pull the valence electrons closer.
Down a group: increases. Each new period adds a shell, so the outermost electrons are further from the nucleus.
Cations (metals lose electrons) are smaller than their parent atoms because they lose an entire energy level.
Anions (nonmetals gain electrons) are larger than their parent atoms because the extra electron increases electron-electron repulsion.
Across a period among cations, size decreases (more protons, same number of shells). When you cross from metals to nonmetals, ionic radius jumps dramatically because anions retain their shells and gain electrons.
Down a group, ionic radius increases (more shells).
The energy required to remove the outermost electron from a gaseous atom.
Across a period: increases. Higher nuclear charge holds electrons more tightly.
Down a group: decreases. Electrons are further from the nucleus and easier to remove.
Ionization energy breaks: half-filled or fully filled subshells (s², p³, p⁶) are extra stable. It takes more energy to remove an electron from these configurations, causing breaks in the otherwise smooth trend.
Inversely related to atomic radius: smaller atom = higher IE.
The degree to which an atom attracts electrons in a chemical bond.
Across a period: increases. Down a group: decreases.
Determines bond type: difference < 0.5 = nonpolar covalent, 0.5 to 1.69 = polar covalent, > 1.7 = ionic.
The energy change when a gaseous atom gains an electron.
Across a period: increases (atoms on the right "want" electrons more). Down a group: decreases.
Metals: reactivity increases going down a group (easier to lose electrons) and decreases across a period.
Nonmetals: reactivity decreases going down a group and increases across a period.
Students often think elements in the same group are identical in properties. They are similar, but inner electrons differ and produce real differences in behaviour (lithium is quite different from caesium, for instance).
Students sometimes assume that ionization energy increases smoothly across every period. It does not. Half-filled and fully filled subshells create "breaks" where IE is higher than expected for that element and then drops for the next.
The direction of reactivity trends for metals and nonmetals is opposite, and students frequently apply the metal trend to nonmetals. Metals get more reactive going down; nonmetals get less reactive going down.
Cations are not "bigger because they lost something." Losing an electron removes repulsion and often an entire shell, making the cation smaller than its parent atom.
⚠️ Know every trend's direction both across a period and down a group. A table question may ask you to rank atoms or ions by size, IE or electronegativity.
⚠️ Be able to explain why each trend occurs (nuclear charge vs. shielding and distance).
⚠️ Electronegativity difference thresholds (< 0.5 nonpolar, 0.5 to 1.69 polar, > 1.7 ionic) are commonly tested.
⚠️ Know the ionization energy break pattern for half-filled and fully filled subshells.
⚠️ The cation-to-anion jump in ionic radius across a period is a classic exam question.
True or false: Atomic radius increases from left to right across a period. (False, it decreases.)
Fill in the blank: Electronegativity ____ across a period and ____ down a group. (increases, decreases)
True or false: A sodium cation (Na⁺) is larger than a neutral sodium atom. (False, it is smaller.)
Fill in the blank: The most reactive metals are found at the - of the periodic table. (bottom-left)
True or false: Noble gases are highly reactive nonmetals. (False, they are stable and largely unreactive.)
Q: Rank the following in order of increasing atomic radius: Na, Mg, K.
A: Mg < Na < K. Mg is to the right of Na in period 3, so it is smaller. K is below Na in Group 1, so it is larger.
Q: Why is the ionization energy of nitrogen higher than that of oxygen, even though oxygen is further right on the periodic table?
A: Nitrogen has a half-filled 2p subshell (2p³), which is extra stable. Removing an electron from that configuration requires more energy than removing one from oxygen's 2p⁴, where one orbital is already paired.
Q: An electronegativity difference of 1.2 between two bonded atoms produces what type of bond?
A: A polar covalent bond (the difference falls in the 0.5 to 1.69 range).
Q: Explain why chloride (Cl⁻) has a larger ionic radius than a neutral chlorine atom.
A: Cl⁻ has gained an electron, increasing electron-electron repulsion while the nuclear charge stays the same. The electron cloud expands, so the ion is larger.
Q: Why does metal reactivity increase going down Group 1?
A: As you go down the group, the outermost electron is further from the nucleus and more shielded by inner electrons, so it is easier to lose. Alkali metals react by losing that one valence electron.
Periodic trends explain why certain chemical reactions happen: ionization energy and electronegativity determine whether atoms form ionic or covalent bonds. The reactivity trends feed directly into understanding replacement reactions (the activity series). Molecular geometry and hybridisation depend on knowing how many valence electrons an atom contributes, which comes straight from its group number.
periodic table, periodic trends, atomic radius, ionic radius, ionization energy, electronegativity, electron affinity, reactivity, alkali metals, alkaline earth metals, halogens, noble gases, transition metals, lanthanides, actinides, s-block, p-block, d-block, f-block, valence electrons, diagonal relationship, alloy, interstitial alloy, substitutional alloy, ferromagnetic, allotrope, nuclear charge, shielding, effective nuclear charge, cation, anion, ionization energy breaks, half-filled subshell