Chapter 3: Compounds—How Elements Combine

3.1 Electron Arrangements and the Octet Rule

Understanding how electrons are arranged in atoms is fundamental to predicting how elements combine to form compounds. The stability of electron arrangements drives chemical bonding.

• Electron Cloud: The region outside the nucleus where electrons are likely to be found.

• Energy Levels: Electrons occupy specific energy levels, starting with the lowest (closest to the nucleus).

• Maximum Electrons per Energy Level: The formula for the maximum number of electrons in a given energy level is , where n is the energy level number.

• Valence Electrons: Electrons in the highest energy level; these are involved in chemical bonding.

• Octet Rule: Atoms tend to achieve eight electrons in their valence shell for stability, like noble gases.

• Periodic Table Relationships: Elements in the same group have the same number of valence electrons.

Example: Noble gases (Group 8A) are stable because they have eight valence electrons (except helium, which has two).

3.2 In Search of an Octet, Part 1: Ion Formation

Atoms can gain or lose electrons to achieve a stable octet, resulting in the formation of ions.

• Ions: Charged atoms formed by gaining or losing electrons.

• Anions: Negatively charged ions formed by gaining electrons.

• Cations: Positively charged ions formed by losing electrons.

• Isoelectronic: Ions with the same number of electrons as the nearest noble gas.

• Trends: Main group elements in Groups 1A, 2A, and 3A form cations; Groups 5A, 6A, and 7A form anions.

• Polyatomic Ions: Groups of nonmetals forming a charged ion; the charge applies to the whole group.

Example: Sodium (Na) loses one electron to form Na+ (cation); chlorine (Cl) gains one electron to form Cl- (anion).

3.3 Ionic Compounds—Electron Give and Take

Ionic compounds are formed when metals and nonmetals combine, transferring electrons to form oppositely charged ions.

• Ionic Bond: The electrostatic attraction between cations and anions.

• Formula: Ionic compounds are written so the total charge is zero; subscripts indicate the number of each ion.

• Naming: Combine the names of the ions, cation first; omit the word "ion." Use Roman numerals for transition metals.

Example: NaCl is sodium chloride; CuO is copper(II) oxide.

3.4 In Search of an Octet, Part 2: Covalent Bond Formation

Nonmetals achieve an octet by sharing electrons, forming covalent bonds and covalent compounds.

• Covalent Bond: A bond formed by sharing valence electrons between atoms.

• Molecule: The smallest unit of a covalent compound.

• Electron-Dot Symbols: Visual representation of valence electrons (Lewis structures).

• Bonding: The number of covalent bonds equals the number of electrons needed to complete the octet.

• Carbon: Forms four bonds (single, double, or triple) to complete its octet.

• Naming: Binary covalent compounds use prefixes to indicate the number of each element; the second element ends in -ide.

Example: Methane (CH4) is a covalent compound; water (H2O) is named by its traditional name.

3.5 The Mole: Counting Atoms and Compounds

The mole is a counting unit used to relate the mass of substances to the number of atoms or molecules present.

• Mole: 1 mole = particles (Avogadro's number).

• Molar Mass: The mass of one mole of a substance, numerically equal to its atomic or formula mass in grams.

• Formula Weight: The sum of atomic masses in a compound.

• Conversions: Use Avogadro's number and molar mass to convert between grams, moles, and number of particles.

Example: 1 mole of carbon-12 weighs 12 grams and contains atoms.

3.6 Getting Covalent Compounds into Shape

The shape of covalent compounds is determined by the arrangement of electron pairs around the central atom, using the VSEPR model.

• Lewis Structure: Shows the arrangement of atoms and electrons.

• VSEPR Model: Valence-Shell Electron-Pair Repulsion model predicts molecular geometry based on repulsion between electron pairs.

• Bond Angles: Electrons arrange themselves to minimize repulsion, resulting in specific shapes (tetrahedral, trigonal planar, linear).

• Nonbonded Electrons: Affect molecular shape by taking up space and pushing bonded pairs closer together.

Example: Methane (CH4) is tetrahedral; water (H2O) is bent due to lone pairs on oxygen.

3.7 Electronegativity and Molecular Polarity

Electronegativity describes an atom's ability to attract electrons in a bond, influencing whether bonds and molecules are polar or nonpolar.

• Electronegativity: Increases toward fluorine on the periodic table.

• Polar Covalent Bond: Electrons are shared unequally; results in partial charges (δ+ and δ-).

• Nonpolar Covalent Bond: Electrons are shared equally.

• Ionic Bond: Formed when the electronegativity difference is 1.8 or greater.

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

Example: Water is a polar molecule due to its bent shape; carbon dioxide is nonpolar because its linear shape causes dipoles to cancel.

Summary Table: Types of Chemical Bonds

Additional info:

• Polyatomic ions often end in -ate or -ite, but some exceptions exist (hydroxide, hydronium, cyanide).

• Dalton (Da) is used in biochemistry for large molecules, equivalent to atomic mass unit (amu).