Chemical Bonding, Naming Compounds, and VSEPR Shapes

Ionic Bonds: Electron Transfer

Ionic bonds form when electrons are transferred between atoms, resulting in positively charged cations (+) and negatively charged anions (-).

Covalent Bonds: Electron Sharing

Covalent bonds form when atoms share electrons. This can be visualized by connecting the dots in Lewis structures.

Electronegativity Difference (|ΔEneg|) and Bond Type

• 2.1 - 3.3: Ionic
• 1.7 - 2.0: Borderline (use other information, e.g., boiling/melting points)
• 0.4 - 1.6: Polar Covalent
• 0.0 - 0.3: Non-Polar Covalent

Drawing Covalent Bonds (Lewis Structures)

1. Place the least electronegative element (not Hydrogen) in the center. Count the total number of valence electrons for all atoms.
2. Connect other atoms to the central atom with single bonds (one shared pair of electrons).
3. Add non-bonding pairs of electrons (lone pairs) to the outer atoms to satisfy the octet rule (usually 8 electrons).
4. Place any remaining extra electrons on the central atom.
5. If the central atom has fewer than an octet, move a non-bonding pair from an exterior atom to form a multiple bond (double or triple) with the central atom.

Drawing Ionic Bonds

Draw the resulting ions (cations and anions) with their respective charges next to each other. Represent the compound by showing the simplest ratio of ions needed for the total charge to sum to zero (the formula unit).

Polyatomic Ions

A polyatomic ion is a group of covalently bonded atoms that collectively carry a net positive or negative charge because the group has gained or lost electrons.

Resonance Structures

Resonance occurs when a single Lewis structure cannot adequately represent a molecule or ion. Multiple valid Lewis structures (resonance structures) are needed, and the actual structure is an average or hybrid of these. You must draw all valid resonance structures.

Ionic Compounds: Properties and Formulas

• The chemical formula represents a Formula Unit, indicating the simplest whole-number ratio of ions required to achieve charge neutrality.
• Ionic compounds typically form crystal lattices.
• They generally have high melting and boiling points.
• They are often brittle solids.

Covalent Naming Prefixes

• 1: Mono-
• 2: Di-
• 3: Tri-
• 4: Tetra-
• 5: Penta-
• 6: Hexa-
• 7: Hepta-
• 8: Octa-
• 9: Nona-
• 10: Deca-
• 11: Undeca-
• 12: Dodeca-

Naming Binary Covalent Compounds

1. Name the first element (usually the least electronegative) using its full element name. Use a prefix (from the list above) to indicate the number of atoms, but omit "mono-" for the first element.
2. Name the second element (more electronegative) by taking its root and adding the suffix "-ide". Use a prefix to indicate the number of atoms, including "mono-".
3. Note: If a prefix ends in "a" or "o" and the element name begins with a vowel, drop the last letter ("a" or "o") from the prefix (e.g., "tetroxide" not "tetraoxide").

Naming Ionic Compounds

1. Identify the cation (positive ion) and the anion (negative ion) in the formula unit.
2. Name the cation using its element name.
3. If the cation can have multiple possible charges (common for transition metals in Groups 3-11 and some elements in Groups 14-16), indicate its positive charge using a Roman numeral in parentheses immediately after the element name (e.g., Iron(II), Lead(IV)).
4. Name the monatomic anion by taking the element root and adding the suffix "-ide" (e.g., Chloride, Oxide).
5. If the anion is a polyatomic ion, use its specific name (e.g., Sulfate, Nitrate). Do not use prefixes to indicate the number of ions in ionic compound names.

Naming Binary Acids

Binary acids consist of Hydrogen and one other nonmetal. Naming convention: "Hydro-" + nonmetal root + "-ic Acid".

• HCl: Hydrochloric Acid
• HBr: Hydrobromic Acid
• HI: Hydroiodic Acid
• HF: Hydrofluoric Acid
• HCN: Hydrocyanic Acid (Note: Cyanide (CN⁻) is a polyatomic ion, but HCN is often named like a binary acid).

Naming Oxyacids

Oxyacids contain Hydrogen, Oxygen, and another nonmetal (often forming a polyatomic ion).

• If the polyatomic ion name ends in "-ate", change the ending to "-ic Acid".
• If the polyatomic ion name ends in "-ite", change the ending to "-ous Acid".

Examples:

• HNO₃ (from Nitrate, NO₃^-): Nitric Acid
• HNO₂ (from Nitrite, NO₂^-): Nitrous Acid
• H₂SO₄ (from Sulfate, SO₄^2-): Sulfuric Acid
• H₂SO₃ (from Sulfite, SO₃^2-): Sulfurous Acid
• H₂CO₃ (from Carbonate, CO₃^2-): Carbonic Acid
• H₃PO₄ (from Phosphate, PO₄^3-): Phosphoric Acid
• CH₃COOH: Acetic Acid (Common name, an organic acid)

VSEPR Theory: Molecular Geometry

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the 3D shape of molecules based on minimizing repulsion between electron pairs around a central atom.

• A = Central atom
• B = Atom(s) bonded to the central atom
• E = Lone pair(s) of electrons on the central atom

Common Geometries (Formula: Electron Geometry, Molecular Geometry, Hybridization)

• AB₂: Linear, Linear, sp
• AB₃: Trigonal Planar, Trigonal Planar, sp²
• AB₄: Tetrahedral, Tetrahedral, sp³
• AB₂E: Trigonal Planar, Bent, sp²
• AB₃E: Tetrahedral, Trigonal Pyramidal, sp³
• AB₂E₂: Tetrahedral, Bent, sp³
• AB₅: Trigonal Bipyramidal, Trigonal Bipyramidal, sp³d
• AB₄E: Trigonal Bipyramidal, Seesaw, sp³d
• AB₃E₂: Trigonal Bipyramidal, T-shaped, sp³d
• AB₂E₃: Trigonal Bipyramidal, Linear, sp³d
• AB₆: Octahedral, Octahedral, sp³d²
• AB₅E: Octahedral, Square Pyramidal, sp³d²
• AB₄E₂: Octahedral, Square Planar, sp³d²