Chemical Bonding

Sharing vs Transfer of Electrons

Chemical bonds form through the interaction of electrons between atoms. The nature of this interaction determines the type of bond.

• Sharing of electrons: Occurs in covalent bonds, where atoms share pairs of electrons.

• Transfer of electrons: Occurs in ionic bonds, where one atom donates electrons to another, resulting in ions.

• Example: Sodium chloride (NaCl) forms by transfer; water (H2O) forms by sharing.

Ionic vs Covalent Compounds

Ionic and covalent compounds differ in their bonding and properties.

• Ionic compounds: Formed from metals and nonmetals; consist of positive and negative ions.

• Covalent compounds: Formed from nonmetals; consist of molecules with shared electrons.

• Example: NaCl (ionic), CO2 (covalent).

Core vs Valence Electrons

Electrons in atoms are classified based on their energy levels.

• Core electrons: Inner electrons not involved in bonding.

• Valence electrons: Outermost electrons involved in chemical reactions and bonding.

• Example: In sodium (Na), 1 valence electron, 10 core electrons.

Octet Rule

The octet rule states that atoms tend to gain, lose, or share electrons to achieve eight valence electrons.

• Key Point: Stable electron configuration similar to noble gases.

• Example: Oxygen forms two bonds to complete its octet.

Single vs Double vs Triple Bonds

Covalent bonds can involve one, two, or three pairs of shared electrons.

• Single bond: One pair of electrons shared (e.g., H–H).

• Double bond: Two pairs shared (e.g., O=O).

• Triple bond: Three pairs shared (e.g., N≡N).

Dipole Moment

A dipole moment arises when there is a separation of charge in a molecule due to differences in electronegativity.

• Key Point: Indicates molecular polarity.

• Example: Water (H2O) has a strong dipole moment.

Electron and Molecular Geometries

The arrangement of atoms and electron pairs determines molecular shape.

• Electron geometry: Based on all electron groups.

• Molecular geometry: Based on only atoms (ignoring lone pairs).

• Example: Methane (CH4) is tetrahedral.

Gases

All Gas Laws

Gas laws describe the relationships between pressure, volume, temperature, and amount of gas.

• Boyle's Law: (at constant T and n)

• Charles's Law: (at constant P and n)

• Avogadro's Law: (at constant P and T)

• Ideal Gas Law:

Intermolecular Forces and Properties

All Intermolecular Attractive Forces

Intermolecular forces are attractions between molecules, affecting physical properties.

• Types: London dispersion, dipole-dipole, hydrogen bonding, ion-dipole.

• Example: Water exhibits hydrogen bonding.

Polarizability

Polarizability is the ability of an electron cloud to be distorted, influencing intermolecular forces.

• Key Point: Larger atoms/molecules are more polarizable.

Polar vs Nonpolar

Molecules can be polar (unequal sharing of electrons) or nonpolar (equal sharing).

• Polar: Water (H2O)

• Nonpolar: Methane (CH4)

Hydrophilic vs Hydrophobic

Hydrophilic substances interact well with water; hydrophobic substances do not.

• Hydrophilic: Polar or ionic compounds.

• Hydrophobic: Nonpolar compounds.

Solutions and Solubility

Homogenous Solutions

A homogenous solution is a mixture with uniform composition throughout.

• Example: Salt water.

Formula and Experimental Mass Analysis

Mass analysis involves determining the composition of compounds.

• Formula mass: Sum of atomic masses in a formula.

• Experimental mass: Determined by measurement.

Insoluble vs Soluble

Solubility describes whether a substance dissolves in a solvent.

• Soluble: Dissolves easily (e.g., NaCl in water).

• Insoluble: Does not dissolve (e.g., AgCl in water).

Viscosity

Viscosity is a measure of a fluid's resistance to flow.

• High viscosity: Honey

• Low viscosity: Water

Polyatomic Ions and Acids

Polyatomic Ions

Polyatomic ions are charged species composed of multiple atoms.

• Example: Nitrate (), sulfate ()

Oxy Acids

Oxy acids contain hydrogen, oxygen, and another element.

• Example: Sulfuric acid ()

Redox and Acid-Base Reactions

Oxidation/Reduction

Redox reactions involve the transfer of electrons.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Example:

Neutralization Reactions

Neutralization occurs when an acid reacts with a base to form water and a salt.

• Example:

Phase Changes and Properties

Evaporation vs Condensation

Evaporation is the process of liquid turning into gas; condensation is gas turning into liquid.

• Evaporation: Endothermic process.

• Condensation: Exothermic process.

Chemical Formulas and Reactions

Empirical vs Molecular Formula

The empirical formula shows the simplest ratio of elements; the molecular formula shows the actual number of atoms.

• Example: Glucose: Empirical (CH2O), Molecular (C6H12O6)

Decomposition

Decomposition reactions break down a compound into simpler substances.

• Example:

Combustion

Combustion reactions involve a substance reacting with oxygen to produce energy.

• Example:

Atomic Structure

Orbitals

Orbitals are regions in an atom where electrons are likely to be found.

• Types: s, p, d, f

Electron Configuration

Electron configuration describes the arrangement of electrons in an atom.

• Example: Sodium:

Orbital Diagram

Orbital diagrams use boxes and arrows to show electron placement and spin.

• Key Point: Follows the Pauli exclusion principle and Hund's rule.

Acid-Base Chemistry

Acid-Base Reactions

Acid-base reactions involve the transfer of protons (H+).

• Example:

Conjugate Pairs

Conjugate pairs consist of an acid and its corresponding base after donating or accepting a proton.

• Example: (acid) and (conjugate base)

Thermochemistry

Thermochemistry

Thermochemistry studies energy changes in chemical reactions.

• Key Point: Endothermic (absorbs energy), exothermic (releases energy).

• Equation:

Lewis Structures

Lewis Dot Structures

Lewis dot structures represent valence electrons as dots around element symbols.

• Key Point: Used to predict bonding and molecular structure.

• Example: Water: H:O:H

Additional info: Academic context and examples have been added to ensure completeness and clarity for exam preparation.