Periodicity and Periodic Trends in the Periodic Table

Periodicity and the Periodic Table

Modern Periodic Table

The modern periodic table is organized by increasing atomic number (Z) from left to right across each row. This arrangement reflects recurring trends in the properties of elements, known as periodicity.

Groups vs Periods

• Group: A vertical column in the periodic table. All elements in a group have the same number of valence electrons, which determines their chemical properties.

• Period: A horizontal row in the periodic table. All elements in a period have the same number of electron shells (principal energy levels).

Example: The 8th element (Oxygen, Z = 8) has the electron configuration 2, 6. It is in Group 6 and the 2nd Period, and forms -2 ions.

Practice: Apply these principles to fluorine (9F) and calcium (20Ca) to determine their positions in the table.

Classification of Elements

Metals, Non-metals, and Metalloids

Type	Location on PT	Elements	Chemical Properties	Physical Properties
Metal	Left of stepped diagonal line	Group 1, Group 2, left of stepped diagonal in Group 3-18, Lanthanoids, Actinoids	Easy to lose valence electron, be oxidized, form cation	Good conductor, malleable, ductile, lustrous
Non-metal	Right of stepped diagonal line	H, right of stepped diagonal in Group 3-18	Easy to gain electron, be reduced, form anion	Poor conductor, non-malleable
Metalloid	Next to stepped diagonal line	B, Si, Ge, As, Sb, Te, At	Both	Semi-conductor

Main Group, Transition Elements, and s, p, d, f Blocks

• Main Group: Group 1 (except H), Group 2, and Groups 13-18 (s-block and p-block elements).

• Transition Elements: Groups 3-11 (d-block elements).

• Lanthanoids: Elements from La (57) to Lu (71).

• Actinoids: Elements from Ac (89) to Lr (103).

• f-block: Sometimes called "Inner Transition Elements".

Periodic Trends

1. Atomic Radius

• Bonding Atomic Radius (Rb): Half the distance between the nuclei of two atoms bonded together in a diatomic molecule.

• Non-bonding Atomic Radius (Rnb): Also called van der Waals radius; relevant for noble gases and non-bonded atoms.

Factors Influencing Atomic Radius:

1. Energy Level: More electron shells mean electrons are further from the nucleus (increases atomic radius).

2. Charge on Nucleus: Higher nuclear charge pulls electrons closer (decreases atomic radius).

3. Shielding Effect: Inner electrons repel outer electrons, increasing atomic radius.

Trends:

• Across a Period: Atomic radius decreases from left to right due to increasing nuclear charge and similar shielding.

• Down a Group: Atomic radius increases due to additional electron shells and increased shielding.

2. Ionic Radius

• Isoelectronic Ions: Ions with the same number of electrons (e.g., Al3+, Mg2+, Na+, Ne, F-, O2-, N3- all have 10 electrons). More protons result in a smaller radius.

• Trends:

  • Down a Group: Ionic radius increases for both cations and anions due to more electron shells and increased shielding.

  • Across a Period: Ionic radius decreases due to increased nuclear charge, but there is a jump between cations and anions.

3. Ionization Energy (IE)

• First Ionization Energy (IE1): Minimum energy required to remove one electron from a neutral gaseous atom. $X_{(g)} \rightarrow X^+_{(g)} + e^-$

• Second Ionization Energy (IE2): Energy required to remove a second electron from the ion. $X^+_{(g)} \rightarrow X^{2+}_{(g)} + e^-$

Factors Affecting IE:

• Greater nuclear charge increases IE.

• Greater distance from nucleus decreases IE.

• Greater shielding effect decreases IE.

Trends:

• Across a Period: IE increases due to higher nuclear charge and similar shielding.

• Down a Group: IE decreases due to increased distance and shielding.

Example: Helium (He) has a greater IE than Hydrogen (H) because they have the same number of shells and shielding, but He has a higher nuclear charge.

4. Electronegativity (EN)

• Definition: The measure of attraction between an atom and its shared pair of electrons in a covalent bond.

• High EN means the atom pulls the electron pair more strongly.

• Non-metals have high EN; metals have low EN.

Trends:

• Across a Period: EN increases.

• Down a Group: EN decreases.

Electronegativity trends are similar to ionization energy but opposite to atomic radius.

5. Metallic vs Non-metallic Character

• Across a Period: Metallic character decreases; non-metallic character increases.

• Down a Group: Metallic character increases (metals become more reactive); non-metallic character decreases.

6. Metal Oxide vs Non-metal Oxide (Oxides of the 3rd Period)

Formula of Oxide	Na2O(s)	MgO(s)	Al2O3(s)	CO2(g)	SiO2(s)	P4O10(s)	SO2(g), SO3(g)
Nature of Oxide	Basic	Basic	Amphoteric	Acidic	Acidic	Acidic	Acidic

• Metal Oxide (Basic Oxide):

  • Metal oxide + H2O → Base $\mathrm{Na_2O(s) + H_2O(l) \rightarrow 2NaOH(aq)}$ $\mathrm{MgO(s) + H_2O(l) \rightarrow Mg(OH)_2(aq)}$

  • Metal oxide + Acid → Salt + H2O $\mathrm{CaO(s) + 2HCl(aq) \rightarrow CaCl_2(aq) + H_2O(l)}$

• Non-metal Oxide (Acidic Oxide):

  • Non-metal oxide + H2O → Acid $\mathrm{CO_2(g) + H_2O(l) \rightarrow H_2CO_3(aq)}$ $\mathrm{SO_3(g) + H_2O(l) \rightarrow H_2SO_4(aq)}$ $\mathrm{P_4O_{10}(s) + 6H_2O(l) \rightarrow 4H_3PO_4(aq)}$

  • Non-metal oxide + Base → Salt + H2O $\mathrm{SO_3(g) + 2NaOH(aq) \rightarrow Na_2SO_4(aq) + H_2O(l)}$

• Amphoteric Oxide: (e.g., Al2O3) Acts as both acid and base.

  • As a base (reacts with acid): $\mathrm{Al_2O_3(s) + 6HCl(aq) \rightarrow 2AlCl_3(aq) + 3H_2O(l)}$

  • As an acid (reacts with base): $\mathrm{Al_2O_3(s) + 2NaOH(aq) + 3H_2O(l) \rightarrow 2NaAl(OH)_4(aq)}$

Note: Amphoteric substances can act as both acids and bases. Amphiprotic substances can both donate and accept H+ ions (a subset of amphoteric).

7. Chemical Reactivity

• Group 1 (Alkali Metals):

  • All have one valence electron and similar reactivity patterns.

  • Reactivity increases down the group as the valence electron is less tightly held.

  • Reactions:

    • With water: Metal hydroxide + hydrogen gas $2\mathrm{Na(s) + 2H_2O(l) \rightarrow 2NaOH(aq) + H_2(g)}$ $2\mathrm{K(s) + 2H_2O(l) \rightarrow 2KOH(aq) + H_2(g)}$

    • With halogens: Metal halide $2\mathrm{Na(s) + Cl_2(g) \rightarrow 2NaCl(aq)}$

• Group 17 (Halogens):

  • Halogens higher up in the group are more reactive (easier to gain electrons, better oxidizing agents).

  • Displacement reactions:

    • $\mathrm{Cl_2(g) + 2KI(aq) \rightarrow I_2(s) + 2KCl(aq)}$

    • $\mathrm{Cl_2(g) + 2NaBr(aq) \rightarrow Br_2(l) + 2NaCl(aq)}$

    • $\mathrm{Br_2(g) + 2LiI(aq) \rightarrow I_2(s) + 2LiBr(aq)}$

8. Melting Point (MP)

• Group 1: MP decreases down the group because the strength of the metallic bond decreases (more electron shells). Less energy is required to overcome the metallic bond.

• Group 17: MP increases down the group because intermolecular forces increase (more electrons). More energy is required to overcome these forces.