UNIT 6 - BONDING
I. Ionic Bonding
A. Behavior of metals
1. formation of cations, charges of metal ions
2. size of ions vs. atoms
3. metallic bonding
B. Behavior of nonmetals
1. formation of anions, charges of nonmetal ions
2. size of ions vs. atoms
C. Crystal Formation
1. formula unit
2. empirical formula
D. Properties of ionic compounds
II. Covalent Bonding
A. Diatomic molecules
1. single bonds
2. multiple bonds
B. Molecular compounds
1. molecular formulas
2. structural formulas
3. Lewis structures
C. Shapes of molecules
1. linear and bent
2. trigonal planar and trigonal pyramid
D. Pure and polar covalent compounds
1. electronegativity values
2. determining % ionic character
3. properties of pure and polar covalent compounds
III. Polyatomic Ions
A. Lewis structures
B. Bonding nature in ionic compounds
IV. Intermolecular Forces
A. Hydrogen Bonding
B. Dipole-Dipole Forces
C. London Dispersion Forces
Chm.1.2 Understand the bonding that occurs in simple compounds in terms of bond type, strength, and
1.2.1 Compare (qualitatively) the relative strengths of ionic, covalent, and metallic bonds.
• Describe metallic bonds: “metal ions plus ‘sea’ of mobile electrons”.
• Describe how ions are formed and which arrangements are stable.
• Appropriately use the term cation as a positively charged ion and anion as negatively charged ion.
• Apply the concept that sharing electrons form a covalent compound that is a stable (inert gas) arrangement.
• Draw Lewis (dot diagram) structures for simple compounds and diatomic elements indicating single, double or triple bonds.
1.2.2 Infer the type of bond and chemical formula formed between atoms.
• Determine that a bond is predominately ionic by the location of the atoms on the Periodic Table (metals combined with nonmetals) or when ∆EN > 1.7.
• Determine that a bond is predominately covalent by the location of the atoms on the Periodic Table (nonmetals combined with nonmetals) or when ∆EN < 1.7.
• Predict chemical formulas of compounds using Lewis structures.
1.2.3 Compare inter- and intra- particle forces.
• Explain why intermolecular forces are weaker than ionic, covalent or metallic bonds.
• Explain why hydrogen bonds are stronger than dipole-dipole forces which are stronger than dispersion forces.
• Apply the relationship between bond energy and length of single, double, and triple bonds.
• Describe intermolecular forces for molecular compounds.
‣ H-bond as attraction between molecules when H is bonded to O, N, or F. Dipole-dipole attractions between polar molecules.
‣ London dispersion forces (electrons of one molecule attracted to nucleus of another molecule) – i.e. liquefied inert gases.
‣ Relative strengths (H>dipole>London/van der Waals).
1.2.5 Compare the properties of ionic, covalent, metallic, and network compounds.
• Explain how ionic bonding in compounds determines their characteristics: high MP, high BP, brittle, and high electrical conductivity either in molten state or in aqueous solution.
• Explain how covalent bonding in compounds determines their characteristics: low MP, low BP, poor electrical conductivity, polar nature, etc.
• Explain how metallic bonding determines the characteristics of metals: high MP, high BP, high conductivity, malleability, ductility, and luster.
• Apply Valence Shell Electron Pair Repulsion Theory (VSEPR) for these electron pair geometries and molecular geometries, and bond angles - Electron pair - Molecular (bond angle); Linear framework – linear; Trigonal planar framework– trigonal planar, bent; Tetrahedral framework– tetrahedral, trigonal pyramidal, bent; Bond angles (include distorting effect of lone pair electrons – no specific angles, conceptually only)
• Describe bond polarity. Polar/nonpolar molecules (relate to symmetry) ; relate polarity to solubility—“like dissolves like”