Ionic Compounds

• Giant Ionic Lattice: Ionic compounds organize into a regular, repeating 3D arrangement where each ion is surrounded by ions of the opposite charge to maximize attraction and minimize repulsion.

• Lattice Energy: This is the energy released when gaseous ions combine to form one mole of an ionic solid; it is a measure of the strength of the ionic bonds within the compound.

• Coulombic Forces: The strength of the bond is governed by Coulomb's Law, meaning that ions with higher charges (e.g., $Mg^{2+}$ vs $Na^+$) or smaller ionic radii will form stronger bonds and more stable lattices.

• Coordination Number: This refers to the number of nearest-neighbor ions of opposite charge surrounding a central ion in the lattice, which is determined by the relative sizes of the cations and anions.

Determining Ionic Formulas

• Step 1: Identify Charges: Determine the valency of the metal and non-metal based on their group in the periodic table (e.g., Group 1 is $+1$, Group 17 is $-1$).
• Step 2: Balance Charges: Use the 'criss-cross' method where the numerical value of the charge on one ion becomes the subscript for the other ion.
• Step 3: Simplify: Ensure the subscripts are in the lowest whole-number ratio (e.g., $Mg_2O_2$ must be simplified to $MgO$).

Predicting Solubility

• Ion-Dipole Interactions: To dissolve an ionic compound, water molecules must surround the ions, with the partial charges of water attracting the full charges of the ions.
• Energy Balance: Solubility occurs when the energy released by hydration is sufficient to overcome the lattice energy holding the ions together.

• Charge Neutrality: Always verify that the sum of positive and negative charges in your final formula equals zero.

• State Symbols: In equations, always include $(s)$ for solid ionic compounds and $(aq)$ or $(l)$ when discussing conductivity, as the state determines the property.

• Polyatomic Ions: When using polyatomic ions like $SO_4^{2-}$ or $NO_3^-$, use parentheses if more than one is required in the formula (e.g., $Mg(NO_3)_2$).

• Melting Point Comparisons: If asked to compare melting points, look for the ion with the higher charge first, as charge has a greater impact on lattice strength than ionic radius.

• Solid Conductivity: A common error is stating that ionic solids conduct electricity; they do not, because the ions are locked in fixed positions and cannot move.

• Intermolecular Forces: Students often incorrectly apply the term 'intermolecular forces' to ionic compounds. Ionic compounds are held by 'intramolecular' ionic bonds throughout the entire structure.

• Brittleness Mechanism: Do not confuse brittleness with weakness. Ionic bonds are very strong, but the lattice is rigid; shifting layers causes like-charges to repel, leading to fracture.

• Electronegativity: The degree of 'ionic character' in a bond is a spectrum; a difference in electronegativity greater than 1.7 is generally considered ionic.

• Periodic Trends: Ionic radius decreases across a period and increases down a group, which directly influences the lattice energy and melting points of the resulting compounds.

• Industrial Applications: The high melting points and stability of ionic oxides (refractories) make them essential for lining high-temperature furnaces.