Exam Prep Notes

Ionic Radii

• Positive ions are smaller than their parent atoms due to increased effective nuclear charge.
• Negative ions are larger than their parent atoms due to decreased effective nuclear charge.

Ionic Radius Trends

• Cations and anions generally decrease in size towards the top and right of the periodic table.

Isoelectronic Series

• Ions with the same electronic configuration.
• Increasing nuclear charge in an isoelectronic series leads to a smaller ionic radius.
• Example: S^{2-}, Cl^{-}, Ar, K^{+}$, Ca^{2+}
• As positive charge increases, size decreases because of the greater pull on the same number of electrons.

Ionization Energy (IE)

• Energy required to remove an electron from a gaseous atom.
• Mg(g) \rightarrow Mg^+(g) + e^- \Delta E = 738 kJ/mol = IE_1
• Successive ionization energies increase as electrons are removed.

First Ionization Energy Trends

• Increases from the lower left to the upper right of the periodic table.
• Decreases down a group (electrons are less strongly bound).
• Generally increases across a period (with exceptions).
• Transition metals have relatively similar ionization energies.

Increasing Ionization Energy Across a Period

• Adding a proton and valence electron increases effective nuclear charge (Z_{eff}
• Valence electrons are more strongly attracted, increasing ionization energy.

Dips in Ionization Energy

• Be to B: Starting the 2p subshell makes electron removal easier.
• O vs. N: Paired electrons in oxygen's 2p orbitals are higher in energy, making electron removal easier.

Higher Ionization Energies

• Removing a second electron from sodium requires breaking into its core configuration, resulting in a higher IE than removing a second electron from neon.
• Na \, IE2 = 4560 \, kJ/mol Ne \, IE2 = 3952 \, kJ/mol

Successive Ionization Energies

• Increase smoothly until a core electron is removed, which requires a dramatic jump in energy.

Electron Affinity (EA)

• Energy released when an electron is added to a gaseous atom.
• O(g) + e^- \rightarrow O^-(g) \Delta E = -141 \, kJ/mol = EA
• Increases from lower left to upper right (excluding noble gases).
• Metals have low or no EA; nonmetals have high EA.
• If EA > 0, the electron spontaneously falls off.

Electron Affinities Exceptions

• Nitrogen: EA > 0 due to the need to start pairing electrons.
• Group 2: EA > 0 due to the need to start the p subshell.
• Adding a second electron is never favorable: O^-(g) + e^- \rightarrow O^{2-}(g) \Delta E > 0

Metallic Character

• Metals lose electrons, nonmetals gain them.
• Increased metallic character means lower ionization energy and less favorable electron affinity.
• Most metallic elements are in the bottom left, least metallic in the top right.

Group Properties

• Noble gases: Closed electron shells, very stable.
• Alkali metals: One valence electron, easily removed, form stable ionic solids.
• Halogens: One electron short of a closed shell, large electron affinities, form stable ionic solids.

Bonding Types

• Ionic Bonding: Electron transfer from metal to nonmetal, stabilized by electrostatic attraction.
• Covalent Bonding: Electrons shared between atoms, lowest energy position between nuclei.
• Metallic Bonding: Valence electrons pooled among all metal atoms.

Lewis Symbol

• Represents the nucleus and core electrons; dots represent valence electrons.
• Number of dots = group number.

Lewis Electron Dot Structures

• Represent valence electrons as dots.
• Atoms exchange or share electrons to achieve a stable octet.

Lewis Structures for Ionic Bonds

• Example: 2Na \cdot + S \rightarrow 2Na^+ + S^{2-}

Ionic Compounds

• Formation requires energy to vaporize/ionize the metal and atomize the nonmetal.
• Electron affinity provides some energy back.
• Driving force: electrostatic attraction in the crystal lattice.

Lattice Energy

• Energy emitted when gaseous ions coalesce to form a solid.
• Born-Haber cycle used to calculate lattice energy.

Energy of Electrostatic Attraction

• E = \frac{q1q2}{4 \pi \epsilon_0 r}

Estimating Lattice Energies (LE)

• Compare charges first: largest |q1q2| = largest (negative) LE
• If q1q2 are the same, smaller radius = larger LE