INORGANIC CHEMISTRY — SUMMARISED NOTES
1. History of the Periodic Table
(a) Johann Döbereiner (1829) — Triads
Grouped elements into sets of three (triads) with similar chemical properties. The atomic mass of the middle element was
approximately the average of the other two.
Triad Atomic Masses
Li, Na, K 7, 23, 39
Cl, Br, I 35.5, 80, 127
Ca, Sr, Ba 40, 88, 137
Limitations
● Only a few elements could be arranged into triads.
● Most known elements did not fit the pattern.
(b) John Newlands (1864) — Law of Octaves
Arranged elements in order of increasing atomic mass. Observed that every eighth element had similar properties, like notes
in music.
Example: Li – Be – B – C – N – O – F – Na
Merits
First attempt to show periodicity.
Limitations
● Worked only up to calcium.
● Left no gaps for undiscovered elements.
● Grouped some dissimilar elements together.
(c) Dmitri Mendeleev (1869)
Periodic Law
The physical and chemical properties of elements are periodic functions of their atomic masses.
Features
● Arranged elements in increasing atomic mass.
● Left gaps for undiscovered elements.
● Predicted properties of missing elements.
● Corrected incorrect atomic masses of some elements.
Achievements — predicted:
● Gallium (eka-aluminium)
● Germanium (eka-silicon)
● Scandium (eka-boron)
Limitations
● Position of hydrogen uncertain.
● Isotopes could not be explained.
● Some elements were arranged out of increasing atomic mass.
(d) Henry Moseley (1913)
Modern Periodic Law
,The physical and chemical properties of elements are periodic functions of their atomic numbers. He arranged elements
according to atomic number rather than atomic mass.
Advantages
● Solved isotope problem.
● Correct position of argon and potassium.
● Correct position of cobalt and nickel.
This became the basis of the modern periodic table.
Modern Periodic Table
The modern periodic table contains 7 Periods (horizontal rows) and 18 Groups (vertical columns).
Blocks
● s-block → Groups 1 and 2
● p-block → Groups 13–18
● d-block → Transition elements (Groups 3–12)
● f-block → Lanthanides and Actinides
2. Periodicity
Definition: Periodicity is the regular repetition of physical and chemical properties of elements when arranged in order of
increasing atomic number. This occurs because the valence electron configuration repeats periodically.
3. Periodic Trends in Physical Properties
(A) Atomic Radius
Definition: atomic radius is half the distance between the nuclei of two bonded atoms of the same element.
Trend Across a Period
Atomic radius decreases from left to right.
Reason
● Nuclear charge increases.
● Number of shells remains constant.
● Attraction between nucleus and electrons increases.
● Electrons are pulled closer to the nucleus.
Example: Na > Mg > Al > Si > P > S > Cl
Trend Down a Group
Atomic radius increases down the group.
Reason
● More electron shells are added.
● Increased shielding effect.
● Outer electrons are farther from the nucleus.
Example: Li < Na < K < Rb < Cs
(B) Ionic Radius
Positive Ions (Cations)
Cations are smaller than their atoms.
Example: Na > Na⁺
, Reason: loss of outer shell; greater nuclear attraction on remaining electrons.
Negative Ions (Anions)
Anions are larger than their atoms.
Example: Cl < Cl⁻
Reason: electron-electron repulsion increases; outer shell expands.
(C) Ionisation Energy
Definition: the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of
gaseous positive ions.
Example: Na(g) → Na⁺(g) + e⁻
Trend Across a Period
Ionisation energy increases.
Reason
● Increasing nuclear charge.
● Decreasing atomic radius.
● Stronger attraction for electrons.
Trend Down a Group
Ionisation energy decreases.
Reason
● Larger atomic radius.
● Increased shielding.
● Outer electron is easier to remove.
(D) Electron Affinity
Definition: the energy released when one mole of gaseous atoms gains one mole of electrons.
Example: Cl(g) + e⁻ → Cl⁻(g)
Trend Across a Period
Generally becomes more negative (greater electron affinity).
Reason
● Increasing nuclear charge.
● Greater attraction for incoming electron.
Trend Down a Group
Generally becomes less negative.
Reason
● Larger atomic size.
● Greater shielding.
(E) Electronegativity
Definition: the ability of an atom to attract shared electrons in a covalent bond.
Trend Across a Period
Electronegativity increases.
Reason
● Increased nuclear charge.
1. History of the Periodic Table
(a) Johann Döbereiner (1829) — Triads
Grouped elements into sets of three (triads) with similar chemical properties. The atomic mass of the middle element was
approximately the average of the other two.
Triad Atomic Masses
Li, Na, K 7, 23, 39
Cl, Br, I 35.5, 80, 127
Ca, Sr, Ba 40, 88, 137
Limitations
● Only a few elements could be arranged into triads.
● Most known elements did not fit the pattern.
(b) John Newlands (1864) — Law of Octaves
Arranged elements in order of increasing atomic mass. Observed that every eighth element had similar properties, like notes
in music.
Example: Li – Be – B – C – N – O – F – Na
Merits
First attempt to show periodicity.
Limitations
● Worked only up to calcium.
● Left no gaps for undiscovered elements.
● Grouped some dissimilar elements together.
(c) Dmitri Mendeleev (1869)
Periodic Law
The physical and chemical properties of elements are periodic functions of their atomic masses.
Features
● Arranged elements in increasing atomic mass.
● Left gaps for undiscovered elements.
● Predicted properties of missing elements.
● Corrected incorrect atomic masses of some elements.
Achievements — predicted:
● Gallium (eka-aluminium)
● Germanium (eka-silicon)
● Scandium (eka-boron)
Limitations
● Position of hydrogen uncertain.
● Isotopes could not be explained.
● Some elements were arranged out of increasing atomic mass.
(d) Henry Moseley (1913)
Modern Periodic Law
,The physical and chemical properties of elements are periodic functions of their atomic numbers. He arranged elements
according to atomic number rather than atomic mass.
Advantages
● Solved isotope problem.
● Correct position of argon and potassium.
● Correct position of cobalt and nickel.
This became the basis of the modern periodic table.
Modern Periodic Table
The modern periodic table contains 7 Periods (horizontal rows) and 18 Groups (vertical columns).
Blocks
● s-block → Groups 1 and 2
● p-block → Groups 13–18
● d-block → Transition elements (Groups 3–12)
● f-block → Lanthanides and Actinides
2. Periodicity
Definition: Periodicity is the regular repetition of physical and chemical properties of elements when arranged in order of
increasing atomic number. This occurs because the valence electron configuration repeats periodically.
3. Periodic Trends in Physical Properties
(A) Atomic Radius
Definition: atomic radius is half the distance between the nuclei of two bonded atoms of the same element.
Trend Across a Period
Atomic radius decreases from left to right.
Reason
● Nuclear charge increases.
● Number of shells remains constant.
● Attraction between nucleus and electrons increases.
● Electrons are pulled closer to the nucleus.
Example: Na > Mg > Al > Si > P > S > Cl
Trend Down a Group
Atomic radius increases down the group.
Reason
● More electron shells are added.
● Increased shielding effect.
● Outer electrons are farther from the nucleus.
Example: Li < Na < K < Rb < Cs
(B) Ionic Radius
Positive Ions (Cations)
Cations are smaller than their atoms.
Example: Na > Na⁺
, Reason: loss of outer shell; greater nuclear attraction on remaining electrons.
Negative Ions (Anions)
Anions are larger than their atoms.
Example: Cl < Cl⁻
Reason: electron-electron repulsion increases; outer shell expands.
(C) Ionisation Energy
Definition: the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of
gaseous positive ions.
Example: Na(g) → Na⁺(g) + e⁻
Trend Across a Period
Ionisation energy increases.
Reason
● Increasing nuclear charge.
● Decreasing atomic radius.
● Stronger attraction for electrons.
Trend Down a Group
Ionisation energy decreases.
Reason
● Larger atomic radius.
● Increased shielding.
● Outer electron is easier to remove.
(D) Electron Affinity
Definition: the energy released when one mole of gaseous atoms gains one mole of electrons.
Example: Cl(g) + e⁻ → Cl⁻(g)
Trend Across a Period
Generally becomes more negative (greater electron affinity).
Reason
● Increasing nuclear charge.
● Greater attraction for incoming electron.
Trend Down a Group
Generally becomes less negative.
Reason
● Larger atomic size.
● Greater shielding.
(E) Electronegativity
Definition: the ability of an atom to attract shared electrons in a covalent bond.
Trend Across a Period
Electronegativity increases.
Reason
● Increased nuclear charge.
