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
3. tetrahedral
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
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North Carolina Essential Standards and Clarifying
Objectives:
Chm.1.2 Understand the bonding that occurs in simple compounds in
terms of bond type, strength, and
properties.
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”