Periodic Trends in Properties — Revision Notes

⚡ 30-Second Revision

Atomic Radius: — $\downarrow$ Group (increases), $\rightarrow$ Period (decreases). \n- Ionic Radius: Cations < Parent atom; Anions > Parent atom. Isoelectronic: Size $\propto 1/Z$. \n- Ionization Enthalpy (IE): $\downarrow$ Group (decreases), $\rightarrow$ Period (increases). Exceptions: G13 < G2, G16 < G15. \n- Electron Gain Enthalpy (EGE): $\downarrow$ Group (less negative), $\rightarrow$ Period (more negative). Exceptions: Noble gases (positive), G2/G15 (positive/near zero), Cl > F (more negative). \n- Electronegativity: $\downarrow$ Group (decreases), $\rightarrow$ Period (increases). F is highest. \n- Metallic Character: $\downarrow$ Group (increases), $\rightarrow$ Period (decreases). \n- Non-metallic Character: $\downarrow$ Group (decreases), $\rightarrow$ Period (increases). \n- Oxide Nature: Basic (metals) $\rightarrow$ Amphoteric (metalloids) $\rightarrow$ Acidic (non-metals) across period.

2-Minute Revision

Periodic trends are predictable patterns in elemental properties driven by effective nuclear charge ($Z_{\text{eff}}$), shielding, and electron shells. Atomic and ionic radii generally decrease across a period due to increasing $Z_{\text{eff}}$ and increase down a group due to added shells.

Cations are smaller than parent atoms, anions are larger. For isoelectronic species, size is inversely proportional to nuclear charge. Ionization enthalpy (IE), the energy to remove an electron, generally increases across a period and decreases down a group.

Key exceptions exist: Group 13 elements have lower IE than Group 2, and Group 16 elements have lower IE than Group 15, due to electron configuration stability. Electron gain enthalpy (EGE), the energy change upon adding an electron, typically becomes more negative across a period and less negative down a group.

Notable exceptions include noble gases (positive EGE) and the fact that chlorine has a more negative EGE than fluorine due to reduced electron-electron repulsion in its larger shell. Electronegativity, the ability to attract shared electrons, increases across a period and decreases down a group, with fluorine being the most electronegative.

Metallic character decreases across a period and increases down a group, while non-metallic character shows the opposite trend. The nature of oxides shifts from basic (metals) to amphoteric (metalloids) to acidic (non-metals) across a period.

5-Minute Revision

Periodic trends are systematic variations in physical and chemical properties of elements, fundamentally governed by atomic structure. The three main factors are: **effective nuclear charge ($Z_{\text{eff}}$), which is the net positive charge felt by valence electrons; shielding effect, where inner electrons reduce the nuclear attraction on outer electrons; and the number of electron shells**.

\n\n1. Atomic Radius: Decreases across a period (due to increasing $Z_{\text{eff}}$ pulling electrons closer) and increases down a group (due to addition of new, larger electron shells). \n2. Ionic Radius: Cations are always smaller than their parent atoms, and anions are always larger.

For isoelectronic species (same number of electrons), the ionic radius decreases as the nuclear charge (Z) increases. Example: O$^{2-}$ > F$^-$ > Na$^+$ > Mg$^{2+}$ (all 10 electrons). \n3. Ionization Enthalpy (IE): Energy to remove an electron.

Generally increases across a period (stronger nuclear pull) and decreases down a group (larger size, weaker pull). \n * Exceptions: IE$_1$ of Group 13 < Group 2 (e.g., Al < Mg) because removing a p-electron is easier than breaking a stable s$^2$ configuration.

IE$_1$ of Group 16 < Group 15 (e.g., O < N) because removing an electron from a p$^4$ configuration (O) leads to a stable half-filled p$^3$ configuration, which is favorable. \n4. Electron Gain Enthalpy (EGE): Energy change when an electron is added.

Generally becomes more negative (exothermic) across a period and less negative (less exothermic) down a group. \n * Exceptions: Noble gases, Group 2, and Group 15 elements have positive (endothermic) EGE due to stable electron configurations.

Crucially, EGE of Cl is more negative than F due to the smaller size of F leading to greater electron-electron repulsion in its compact 2p subshell. \n5. Electronegativity: Tendency to attract shared electrons in a bond.

Increases across a period and decreases down a group. Fluorine is the most electronegative element. \n6. Metallic/Non-metallic Character: Metallic character (tendency to lose electrons) increases down a group and decreases across a period.

Non-metallic character (tendency to gain electrons) shows the opposite trend. \n7. Nature of Oxides: Metallic oxides are basic (e.g., Na$_2$O), non-metallic oxides are acidic (e.g., SO$_2$), and some elements near the metal-nonmetal boundary form amphoteric oxides (e.

g., Al$_2$O$_3$, ZnO). Acidity increases across a period, basicity decreases.

Prelims Revision Notes

1
Atomic Radius: — \n * Period: Decreases (L to R) due to increasing $Z_{\text{eff}}$. \n * Group: Increases (T to B) due to increasing 'n' (new shells). \n * Transition elements: Less regular decrease due to d-orbital shielding. \n * Lanthanide Contraction: Poor shielding by 4f electrons leads to smaller than expected radii for 5d series elements (e.g., Zr and Hf have similar sizes). \n2. Ionic Radius: \n * Cation: Smaller than parent atom (loss of shell, increased $Z_{\text{eff}}$). \n * Anion: Larger than parent atom (electron repulsion, decreased $Z_{\text{eff}}$ per electron). \n * Isoelectronic Species: For same number of electrons, size $\propto 1/Z$. Example: O$^{2-}$ > F$^-$ > Na$^+$ > Mg$^{2+}$. \n3. Ionization Enthalpy (IE): Energy to remove 1st electron (IE$_1$). \n * Period: Increases (L to R) due to increasing $Z_{\text{eff}}$ and decreasing size. \n * Group: Decreases (T to B) due to increasing size and shielding. \n * Exceptions: \n * IE$_1$ (Group 13) < IE$_1$ (Group 2) (e.g., Al < Mg) due to ease of removing p-electron vs. stable s$^2$. \n * IE$_1$ (Group 16) < IE$_1$ (Group 15) (e.g., O < N) due to stability of half-filled p$^3$ after removing one electron from p$^4$. \n * IE$_1 <$ IE$_2 <$ IE$_3$ (harder to remove from positive ion). \n4. Electron Gain Enthalpy (EGE): Energy change when adding 1 electron. \n * Period: More negative (L to R, more exothermic) due to increasing $Z_{\text{eff}}$. \n * Group: Less negative (T to B, less exothermic) due to increasing size. \n * Exceptions: \n * Noble gases, Group 2, Group 15: Positive EGE (endothermic) due to stable configurations. \n * Cl has more negative EGE than F (due to smaller size of F causing electron-electron repulsion). \n5. Electronegativity: Tendency to attract shared electrons in a bond. \n * Period: Increases (L to R). \n * Group: Decreases (T to B). \n * Most Electronegative: Fluorine (F). \n6. Metallic Character: Tendency to lose electrons. \n * Period: Decreases (L to R). \n * Group: Increases (T to B). \n7. Non-metallic Character: Tendency to gain electrons. \n * Period: Increases (L to R). \n * Group: Decreases (T to B). \n8. Nature of Oxides: \n * Basic: Formed by metals (e.g., Na$_2$O, CaO). \n * Amphoteric: Formed by elements near metal-nonmetal boundary (e.g., Al$_2$O$_3$, ZnO, PbO, SnO$_2$). \n * Acidic: Formed by non-metals (e.g., CO$_2$, SO$_3$, Cl$_2$O$_7$). \n * Trend: Basic $\rightarrow$ Amphoteric $\rightarrow$ Acidic across a period.

Vyyuha Quick Recall

All Ions Except Noble Gases Exhibit Regular Trends: \n\nAtomic Radius: Increases Down, Decreases Across. \nIonization Energy: Decreases Down, Increases Across.

\nElectron Gain Enthalpy: Less Negative Down, More Negative Across. \nNegativity (Electronegativity): Decreases Down, Increases Across. \n\n*Remember F is the most electronegative, and Cl has more negative EGE than F!