Ionic and Covalent Compounds, Molecular Geometry, and Intermolecular Forces: Study Guide

Study Guide - Smart Notes

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Chapter 5: Ionic Compounds

Ionic Bonding and Ion Formation

Ionic bonding involves the transfer of electrons from one atom to another, resulting in the formation of oppositely charged ions that are held together by electrostatic attraction. This process typically occurs between metals and nonmetals.

Anion Formation: Nonmetals gain electrons to form negatively charged ions (anions).
Cation Formation: Metals lose electrons to form positively charged ions (cations).
Stability: Ions form to achieve noble gas electron configurations, resulting in stable electronic arrangements.

Example: Sodium (Na) loses one electron to form Na+; chlorine (Cl) gains one electron to form Cl-. They combine to form NaCl.

Formulas and Names of Ionic Compounds

Binary Ionic Compounds: Composed of two elements (metal + nonmetal). Name the cation first, then the anion (with -ide ending).
Transition Metals: May have multiple possible charges. Indicate the charge with Roman numerals in parentheses (e.g., FeCl2 is iron(II) chloride).
Polyatomic Ions: Charged groups of covalently bonded atoms (e.g., SO42-, NH4+). Names and formulas are provided on reference sheets.

Example: Ca(NO3)2 is calcium nitrate.

Chapter 6: Covalent Compounds

Formulas and Names of Binary Molecular Compounds

Covalent (molecular) compounds are formed by the sharing of electrons between nonmetals. Prefixes indicate the number of each atom present.

Prefixes: mono-, di-, tri-, tetra-, penta-, hexa-, etc.
Naming: The first element keeps its name; the second element ends with -ide. Prefixes are used for both elements (except mono- is often omitted for the first element).

Example: CO2 is carbon dioxide; N2O4 is dinitrogen tetroxide.

Counting Valence Electrons and Drawing Lewis Structures

Valence Electrons: Electrons in the outermost shell, involved in bonding.
Bonding Electrons: Shared between atoms (represented as lines in Lewis structures).
Non-bonding Electrons: Lone pairs (represented as dots).
Lewis Structures: Show all valence electrons, both bonding and non-bonding, for each atom in a molecule.

Example: The Lewis structure for water (H2O) shows two single bonds and two lone pairs on oxygen.

Covalent Bonding Patterns for Main Group Elements

Octet Rule: Atoms tend to form bonds to achieve eight valence electrons (except hydrogen, which seeks two).
Common Patterns: Carbon forms four bonds, nitrogen three, oxygen two, and halogens one.

Condensed and Skeletal Structures

Condensed Formula: Shows the arrangement of atoms without drawing all bonds (e.g., CH3CH2OH).
Skeletal Structure: Lines represent carbon-carbon bonds; hydrogens attached to carbons are often omitted for simplicity.

Chapter 7: Molecular Geometry and Intermolecular Forces

Shapes and Bond Angles: VSEPR Theory

The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on the repulsion between electron groups around a central atom.

Electron Groups: Include both bonding pairs (single, double, triple bonds) and lone pairs.
Common Geometries and Bond Angles:

Geometry	Electron Groups	Bond Angle
Linear	2	180°
Trigonal Planar	3	120°
Bent (Trigonal Planar)	3 (1 lone pair)	~120°
Tetrahedral	4	109.5°
Trigonal Pyramidal	4 (1 lone pair)	~107°
Bent (Tetrahedral)	4 (2 lone pairs)	~104.5°

Example: Methane (CH4) is tetrahedral; water (H2O) is bent.

Evaluating Molecular Shape with Multiple Central Atoms

Analyze each central atom separately using VSEPR theory.
Overall molecular shape is a combination of the local geometries.

Electronegativity and Bond Polarity

Electronegativity: The ability of an atom to attract shared electrons in a bond.
Bond Polarity: A bond is polar if there is a significant difference in electronegativity between the two atoms.
Use of Electronegativity Values: The greater the difference, the more polar the bond.

Example: The O-H bond in water is polar because oxygen is more electronegative than hydrogen.

Molecular Polarity

Determined by both bond polarity and molecular geometry.
Nonpolar molecules have symmetrical charge distribution; polar molecules have an uneven distribution.

Example: CO2 is nonpolar (linear, dipoles cancel); H2O is polar (bent, dipoles do not cancel).

Intermolecular Forces (IMFs)

Dispersion Forces (London Forces): Weakest, present in all molecules, due to temporary dipoles.
Dipole-Dipole Forces: Occur between polar molecules.
Hydrogen Bonding: Strongest type of dipole-dipole interaction; occurs when H is bonded to N, O, or F.

Melting and Boiling Point Trends

Stronger intermolecular forces lead to higher melting and boiling points.
Hydrogen bonding > dipole-dipole > dispersion forces.

Example: Water has a high boiling point due to hydrogen bonding.

Solubility and Polarity

"Like dissolves like": Polar molecules dissolve in polar solvents; nonpolar molecules dissolve in nonpolar solvents.
Solubility can be predicted by comparing molecular polarity and the nature of the solvent.

Example: NaCl dissolves in water (both polar); oil does not dissolve in water (nonpolar vs. polar).

Polarity of Large Molecules

Large molecules may have both polar and nonpolar regions.
Overall polarity depends on the balance and arrangement of these regions.

Hydrophobic and Hydrophilic Terms

Hydrophobic: "Water-fearing"; nonpolar substances that do not mix well with water.
Hydrophilic: "Water-loving"; polar or charged substances that interact well with water.

Example: Fatty acids have hydrophobic tails and hydrophilic heads, important in biological membranes.

Key Equations and Concepts

Lewis Structure Electron Count:

Formal Charge:

Electronegativity Difference and Bond Type:

Nonpolar covalent:
Polar covalent:
Ionic: