The Periodic Table is chemistry's most powerful organising tool. It arranges elements in a way that reveals deep patterns in their properties. This chapter traces how the table developed and explains the trends that emerge from it.
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Key Concepts
- Historical Development
- Dobereiner's Triads (1829): Grouped elements in sets of three where the middle element's properties were approximately the average of the other two. E.g., Li–Na–K.
- Newlands' Law of Octaves (1865): When elements are arranged in increasing atomic mass, every eighth element shows properties similar to the first. This worked only up to calcium.
- Mendeleev's Periodic Law (1869): Properties of elements are a periodic function of their atomic masses. He left gaps for undiscovered elements and predicted their properties (e.g., eka-aluminium = Gallium).
- Modern Periodic Law (Moseley, 1913): Properties of elements are a periodic function of their atomic numbers.
- Structure of the Modern Periodic Table
- The modern table has 18 groups (vertical columns) and 7 periods (horizontal rows).
- Groups 1–2: s-block elements (alkali and alkaline earth metals)
- Groups 13–18: p-block elements
- Groups 3–12: d-block elements (transition metals)
- Lanthanoids and Actinoids: f-block elements (inner transition metals)
- Period number = number of electron shells; Group number relates to valence electrons (for main group).
Periodic Trends
1. Atomic Radius
Defined as half the distance between nuclei of two adjacent identical atoms. Generally decreases across a period (increasing nuclear charge pulls electrons closer) and increases down a group (additional electron shells). Note: Anions are larger than parent atoms; cations are smaller.
2. Ionic Radius
For isoelectronic species (same number of electrons), ionic radius decreases with increasing atomic number because more protons pull the same electrons closer.
3. Ionisation Enthalpy (IE)
Energy required to remove the outermost electron from a gaseous atom in its ground state. Increases across a period (stronger attraction); decreases down a group (outer electron farther from nucleus). Exceptions: IE of Be > B (2s2 vs 2p1) and IE of N > O (half-filled 2p is stable).
4. Electron Gain Enthalpy (Electron Affinity)
Energy change when an electron is added to a gaseous atom. Generally becomes more negative across a period (more easily gains electron). Note: Noble gases have large positive values; F has slightly less negative value than Cl due to small size and electron-electron repulsion in F.
5. Electronegativity
The tendency of an atom in a molecule to attract shared electrons towards itself. Increases across a period and decreases down a group. Fluorine (F) is the most electronegative element (Pauling scale: 4.0).
6. Valence
The combining capacity of an element. For main group elements, valence = group number (up to 4) or 18 - group number (for groups 14-18 when the higher valence is used).
- Classification of Elements
- Metals: Left and centre of periodic table; malleable, ductile, conductors.
- Non-metals: Upper right; non-conductors (except graphite).
- Metalloids (Semimetals): Along the staircase line (B, Si, Ge, As, Sb, Te); intermediate properties.
- Noble gases (Group 18): Completely filled shells; chemically inert.
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Worked Examples
Arrange F, Cl, Br, I in order of increasing atomic radius.
Atomic radius increases down a group: F < Cl < Br < I
Which has a higher first ionisation enthalpy: Mg or Al?
Mg ([Ne] 3s2) has a filled 3s subshell, which is more stable than Al ([Ne] 3s2 3p1). So IE1 of Mg > Al.
Arrange Na+, Mg2+, F- in order of increasing ionic radius. All are isoelectronic (10 electrons).
More protons = smaller radius. Na (Z=11) > Mg (Z=12) in size... so ionic radii: Mg2+ < Na+ < F-
Which element in the second period has the highest electronegativity?
Fluorine (F) — it is in the top-right corner of the periodic table and has the highest electronegativity of all elements.
Mendeleev predicted the properties of eka-aluminium. Which element was later discovered to match this prediction?
Gallium (Ga), discovered in 1875, matched Mendeleev's predictions almost exactly.
Why does atomic radius decrease across a period?
Across a period, the atomic number (nuclear charge) increases while electrons are added to the same shell. The increased nuclear charge pulls electrons inward, reducing atomic size.
Identify the block and period of the element with electronic configuration [Ar] 3d10 4s2 4p3 (Z=33).
Valence electrons are in 4p, so this is a p-block element. Highest n = 4, so it is in Period 4. It is Arsenic (As).
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Common mistakes
Students often say ionisation enthalpy always increases smoothly across a period. Remember the two dips: group 13 (p1 is less stable than s2, so IE of Al < Mg) and group 16 (O has lower IE than N because the paired 2p electron in O repels more easily). Also, do NOT confuse electron gain enthalpy (which can be negative = exothermic) with ionisation enthalpy (always endothermic).
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Summary
The modern periodic table is arranged by atomic number. Key trends — atomic radius, ionisation enthalpy, electron gain enthalpy, and electronegativity — all arise from changes in nuclear charge and electron shielding. Knowing these trends allows prediction of chemical behaviour without memorising every element individually.