Chapter 4: Periodic Table and Periodicity of Properties

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Chapter Overview & SLOs

What is the periodic table? The periodic table is an organized arrangement of elements in order of increasing atomic number, with elements having similar chemical properties placed in the same vertical columns (groups).

How is the periodic table organized? Elements are arranged into:

• 7 Periods (horizontal rows): Period number corresponds to the highest energy level ($n$) occupied by electrons.
• 8 Main Groups (vertical columns): Group number for main group elements corresponds to the number of valence electrons (electrons in outermost shell).

What are the different blocks in the periodic table?

• s-block (Groups IA and IIA): Valence electrons occupy s orbitals.
  • Group IA: Alkali metals (Li, Na, K, Rb, Cs, Fr) — highly reactive, form +1 ions.
  • Group IIA: Alkaline earth metals (Be, Mg, Ca, Sr, Ba, Ra) — reactive, form +2 ions.
• p-block (Groups IIIA to VIIIA): Valence electrons occupy p orbitals.
  • Group VIIA: Halogens (F, Cl, Br, I, At) — 'salt-formers', highly reactive nonmetals, form -1 ions.
  • Group VIIIA: Noble gases (He, Ne, Ar, Kr, Xe, Rn) — inert, full valence shell (octet).
• d-block (Transition metals): Valence electrons occupy d orbitals.
• f-block (Inner transition metals): Valence electrons occupy f orbitals (lanthanides and actinides).

How do we locate an element in the periodic table? Use the electronic configuration:

• Period: Determined by the highest principal quantum number ($n$) of occupied shells. Example: Aluminum ($1s^2 2s^2 2p^6 3s^2 3p^1$) → highest $n=3$ → Period 3.
• Group: For main group elements, group number = number of valence electrons. Example: Aluminum has 3 valence electrons ($3s^2 3p^1$) → Group IIIA (Group 13).
• Block: Determined by the orbital type of the valence electrons. Example: Aluminum's valence electrons enter p orbital → p-block.

What is shielding effect? The shielding effect is the reduction in effective nuclear charge on an electron due to repulsion from inner shell electrons.

• Trend down a group: Shielding effect increases because more inner electron shells are added, increasing repulsion and reducing the attraction between nucleus and valence electrons.
• Trend across a period: Shielding effect remains relatively constant because electrons are added to the same shell, with little increase in inner shell shielding.

What is ionization energy? Ionization energy is the minimum energy required to remove the most loosely bound electron from a gaseous atom.

• Trend down a group: Ionization energy decreases because atomic size increases, shielding effect increases, making valence electrons easier to remove.
• Trend across a period: Ionization energy increases because nuclear charge increases, atomic radius decreases, and valence electrons are held more tightly.
• Example: Potassium has lower ionization energy than Sodium because it is larger and has more shielding.

What is electron affinity? Electron affinity is the energy released when an electron is added to a gaseous atom.

• Trend down a group: Electron affinity generally decreases (less energy released) because larger atoms have less attraction for additional electrons.
• Trend across a period: Electron affinity generally increases (more energy released) because increased nuclear charge attracts added electrons more strongly.

What is atomic radius? Atomic radius is the distance from the nucleus to the outermost electron shell.

• Trend down a group: Atomic radius increases because new electron shells are added.
• Trend across a period: Atomic radius decreases because increasing nuclear charge pulls electrons closer to the nucleus.

What determines chemical similarity? Elements in the same group have similar valence shell electronic configurations, which means they exhibit similar chemical properties (e.g., all alkali metals react vigorously with water to form hydroxides and hydrogen gas).

Key periodic trends summary:

| Property | Down a Group | Across a Period (Left to Right) | |----------|--------------|--------------------------------| | Atomic radius | Increases | Decreases | | Shielding effect | Increases | Constant | | Ionization energy | Decreases | Increases | | Electron affinity | Decreases (generally) | Increases (generally) | | Metallic character | Increases | Decreases | | Electronegativity | Decreases | Increases |

These notes are strictly aligned with the Student Learning Outcomes (SLOs) for the FBISE 2026 annual examination.

• How do we locate elements in the periodic table? Use valence shell electronic configurations (like $1s^2 2s^2 2p^6 3s^2 3p^1$ for Aluminum) to pinpoint an element's specific period (highest $n$ value) and group (number of valence electrons), and determine its block based on valence orbital (s or p).
• How trends in atomic properties affect reactivity? Explain why ionization energy decreases down a group (due to increased atomic size and shielding effect) and how increased shielding makes valence electrons easier to remove in larger atoms like Potassium compared to Sodium, making Potassium more reactive.
• How do we categorize elements into blocks? Classify elements into the s-block (Groups IA and IIA: alkali metals and alkaline earth metals) or p-block (Groups IIIA to VIIIA: including halogens Group VIIA and noble gases Group VIIIA) based on whether their valence electrons occupy s or p orbitals.
• How do we predict periodic variations? Analyze how nuclear charge (increasing across a period) and atomic size (decreasing across a period) dictate changes in electron affinity and ionization energy as you move across a period or down a group, using the periodic trends table.

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