Periodic Trends: Atomic Size, Ionization Energy, Electron Affinity, Ionic Size, and Metallic Character

Periodic Trends in Atomic Properties

1. Atomic Size (Atomic Radius)

• Definition: The radius of an atom, representing the size of its electron cloud.

• Trends:

  • Across a Period (left to right): Decreases.

    • Reason: As you move across a period, the number of protons (nuclear charge) increases while electrons are added to the same main energy level. This results in a stronger nuclear pull on the electrons, drawing them closer to the nucleus and decreasing the atomic radius.

  • Down a Group (top to bottom): Increases.

    • Reason: As you move down a group, new main energy levels are added, placing the outermost electrons further from the nucleus. This increased distance and the shielding effect from inner electrons diminish the nuclear pull on the outer electrons, leading to a larger atomic size.

• Examples:

  • Comparing S, O, Po, Te (Group 16): Po > Te > S > O (size increases down the group).

  • Comparing Cl, Mg, Ar, Al (Period 3): Mg > Al > Cl > Ar (size decreases across the period).

  • Comparing K, Li, Fr, H (Group 1): Fr > K > Li > H (size increases down the group).

  • Comparing Xe, Rb, Zn: Rb (Group 1) is largest, then Zn (Group 12), then Xe (Group 18), as size generally decreases across a period.

2. Ionization Energy (I.E.)

• Definition: The minimum energy required to remove one electron from a gaseous atom or ion.

  • $X(g) \to X^+(g) + e^-$

• Trends:

  • Across a Period (left to right): Increases.

    • Reason: Atoms become smaller across a period, and the nuclear pull on the valence electrons becomes stronger. Electrons are held more tightly, so more energy is required to remove them (it's harder to take an electron).

  • Down a Group (top to bottom): Decreases.

    • Reason: Atoms become larger down a group, and the valence electrons are further from the nucleus. Increased shielding from inner electrons also reduces the effective nuclear charge felt by the outer electrons, making them easier to remove (less energy needed).

• Factors influencing I.E.:

  • Smaller atom: Stronger nuclear pull holding electrons tighter, harder to take.

  • Shielded electrons: Electrons further from the nucleus are shielded from the full nuclear charge, making them easier to remove.

• Examples:

  • Hardest to take an electron from (Highest I.E.): O, Li, F, or B?

    • F (F has $p^+ = 9$ electrons being pulled tighter in the same energy level as O and B, and is much smaller than Li).

    • Order: F > O > B > Li

  • Hardest to take an electron from (Highest I.E.): Ag, Cu, or Au?

    • Cu (I.E. decreases down a group, so Cu is smaller and has a higher I.E. than Ag and Au).

    • Order: Cu > Ag > Au

• Successive Ionization Energies (I.E.$<em>1$, I.E.$</em>2$, I.E.$_3$…):

  • Successive ionization energies always increase (I.E1 < I.E2 < I.E_3) because it is always harder to remove an electron from a positively charged ion than from a neutral atom, and even harder from a $+2$ ion, and so on.

  • Significant Jumps in I.E.: Occur when an electron is removed from a stable noble gas electron configuration.

    • Lithium (Li):

      • First I.E. ($Li \to Li^+ + e^-$): $520 \text{ kJ/mol}$. (Removes valence electron)

      • Second I.E. ($Li^+ \to Li^{2+} + e^-$): $7297 \text{ kJ/mol}$. (Removes an electron from a core shell, as $Li^+$ now has a stable electron configuration like He ($1s^2$)). This is a huge jump.

    • Magnesium (Mg):

      • First I.E. ($Mg \to Mg^+ + e^-$): $737.6 \text{ kJ/mol}$.

      • Second I.E. ($Mg^+ \to Mg^{2+} + e^-$): $1450 \text{ kJ/mol}$.

      • Third I.E. ($Mg^{2+} \to Mg^{3+} + e^-$): $7732 \text{ kJ/mol}$ (Removes an electron from a core shell, as $Mg^{2+}$ now has a stable electron configuration like Ne ($1s^2 2s^2 2p^6$)). Also a significant jump.

  • General Rule from I.E. Jumps:

    • Elements in Group 1 will have the highest second ionization energy because removing the first electron makes them isoelectronic with a noble gas.

    • Elements in Group 2 will have the highest third ionization energy because removing the first two electrons makes them isoelectronic with a noble gas.

3. Electron Affinity (E.A.)

• Definition: The energy change that occurs when an electron is added to a gaseous atom.

  • $X(g) + e^- \to X^-(g) + \text{energy}$

  • When energy is released, the E.A. value is negative (exothermic process). A more negative value indicates a higher electron affinity (greater attraction for electrons).

• Trends:

  • Across a Period (left to right): Generally increases (becomes more negative) (same general trend as I.E.).

    • Reason: As nuclear charge increases across a period and atomic size decreases, there is a stronger attraction for an incoming electron, leading to more energy being released.

  • Down a Group (top to bottom): Generally decreases (becomes less negative) (same general trend as I.E.).

    • Reason: As atomic size increases down a group, the incoming electron is added further from the nucleus and experiences more shielding, resulting in a weaker attraction and less energy released.

• Examples:

  • Which element releases the most energy when an electron is added (Highest E.A.)? This means asking for the element with the most negative E.A.

    • Si, C, Sn, Pb (Group 14): C (E.A. generally increases up a group).

    • Br, K, Ni, Se: Br (E.A. generally increases across a period, Br is a halogen and very electronegative).

4. Ionic Size

• Definition: The radius of an ion.

• Comparison of Atom vs. Ion:

  • Anions (negatively charged ions): Formed by gaining electrons. They are always larger than their parent neutral atoms.

    • Example: F vs. F$^-$

      • F: $p^+ = 9, e^- = 9$

      • F$^-$(gain e-): $p^+ = 9, e^- = 10$

      • Reason: The nuclear pull ($p^+$) remains the same, but the addition of an extra electron ($e^-$) increases electron-electron repulsion, causing the electron cloud to expand.

  • Cations (positively charged ions): Formed by losing electrons. They are always smaller than their parent neutral atoms.

    • Example: Ca vs. Ca$^{2+}$

      • Ca: $p^+ = 20, e^- = 20$

      • Ca$^{2+}$(lost e-): $p^+ = 20, e^- = 18$

      • Reason: The loss of electrons reduces electron-electron repulsion. If a full shell of electrons is lost, the cation becomes significantly smaller as the outermost occupied energy level is now closer to the nucleus.

5. Metallic Character

• Definition: A measure of how readily an element can lose electrons (high electropositivity) and exhibit properties typical of metals.

• Properties of Metals: Malleable, ductile, good conductors of heat and electricity, shiny luster.

• Trends:

  • Across a Period (left to right): Decreases.

    • Reason: As you move across, atoms hold onto their electrons more tightly (higher I.E.), making it harder to lose them and reducing metallic properties.

  • Down a Group (top to bottom): Increases.

    • Reason: As you move down, atoms are larger and hold onto their valence electrons less tightly (lower I.E.), making it easier to lose them and increasing metallic properties.

• Most Metallic Element: Francium (Fr) is typically considered the most metallic element.

• Examples:

  • Who has the most metallic properties: Cl, Rb, K, or Ni?

    • Rb (Group 1 alkali metal, further down than K). The order of metallic character would be: Rb > K > Ni > Cl (Cl is a non-metal with very low metallic character).

Summary of Periodic Trends