Study Guide: Ionic and Covalent Compounds in GOB Chemistry

Ionic Compounds

Nature and Formation of Ionic Compounds

Ionic compounds are formed by the transfer of electrons from one atom (typically a metal) to another (typically a nonmetal), resulting in the formation of oppositely charged ions that are held together by electrostatic forces.

• Charge Prediction: The charge of a main group ion can be predicted based on its position in the periodic table. For example, Group 1 elements form +1 ions, Group 2 form +2 ions, Group 17 (halogens) form -1 ions, etc.

• Formation Explanation: Ionic compounds form because atoms achieve stable electron configurations (often noble gas configurations) by gaining or losing electrons.

• Formula Prediction: The formula of an ionic compound is determined by balancing the total positive and negative charges so the compound is electrically neutral.

Naming Ions and Ionic Compounds

• Cations and Anions: Cations are positively charged ions (usually metals), and anions are negatively charged ions (usually nonmetals).

• Main Group Cations: Named by the element name (e.g., sodium ion for Na+).

• Transition Metal Cations: Named with the element name and a Roman numeral indicating the charge (e.g., iron(III) ion for Fe3+).

• Polyatomic Ions: Ions composed of more than one atom (e.g., sulfate SO42−, ammonium NH4+).

• Naming Ionic Compounds: Name the cation first, then the anion. For transition metals, include the charge in Roman numerals. For polyatomic ions, use their standard names.

• Writing Formulas: Use the charges to determine the ratio of ions needed for neutrality. For example, for magnesium chloride: Mg2+ and Cl− combine as MgCl2.

General Properties of Ionic Compounds

• High melting and boiling points

• Conduct electricity when molten or dissolved in water

• Usually solid and crystalline at room temperature

Example: Sodium chloride (NaCl) is an ionic compound formed from Na+ and Cl− ions.

Covalent Compounds

Nature and Formation of Covalent Compounds

Covalent compounds are formed when two nonmetal atoms share electrons to achieve stable electron configurations. This sharing results in the formation of molecules.

• Why They Form: Atoms share electrons to achieve a full valence shell, typically following the octet rule.

• Binary Covalent Compounds: Compounds composed of two different nonmetals (e.g., CO2, H2O).

Lewis Dot Structures

• Visual representations of the valence electrons in molecules.

• Show how atoms are bonded and where lone pairs of electrons are located.

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

Molecular Shape and VSEPR Theory

• VSEPR (Valence Shell Electron Pair Repulsion) Theory: Predicts the shape of molecules based on repulsion between electron pairs around a central atom.

• Common Shapes: Linear, bent, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

Electronegativity and Polarity

• Electronegativity: A measure of an atom's ability to attract shared electrons. Fluorine is the most electronegative element.

• Polarity: A molecule is polar if it has an uneven distribution of electron density, resulting in a dipole moment.

• Bond Polarity: Determined by the difference in electronegativity between bonded atoms.

• Molecular Polarity: Determined by both bond polarity and molecular shape.

Intermolecular Forces

• Types of Intermolecular Forces:

  • Dipole-dipole: Attractions between polar molecules.

  • Hydrogen bonding: A strong type of dipole-dipole interaction involving H bonded to N, O, or F.

  • London dispersion forces: Weak attractions due to temporary dipoles in all molecules, especially significant in nonpolar molecules.

  • Induced dipole: Temporary polarization of a molecule due to the presence of a nearby ion or polar molecule.

• Effect on Physical Properties: Stronger intermolecular forces lead to higher boiling and melting points, and affect solubility.

Solids: Types and Forces

• Molecular Solids: Held together by intermolecular forces (e.g., ice).

• Ionic Solids: Held together by ionic bonds (e.g., NaCl).

• Metallic Solids: Metal atoms held together by a 'sea' of delocalized electrons.

• Network Covalent Solids: Atoms connected by covalent bonds in a continuous network (e.g., diamond).

Example: Table salt (NaCl) is an ionic solid, while diamond is a network covalent solid.

Summary Table: Types of Intermolecular Forces

Type of Force	Occurs Between	Relative Strength	Example
London Dispersion	All molecules (especially nonpolar)	Weakest	O2, N2
Dipole-Dipole	Polar molecules	Intermediate	HCl
Hydrogen Bonding	H bonded to N, O, or F	Strongest (of intermolecular)	H2O, NH3

Additional info: The original notes were in outline form; academic context and examples have been added for clarity and completeness.