Chemical Bonding and Nomenclature: Study Guide for General Chemistry

Chemical Bonding

Ionic Bonding

Ionic bonding occurs when atoms transfer electrons, resulting in the formation of oppositely charged ions that attract each other. This typically happens between metals and nonmetals.

• Definition: The attractive force between a cation (positively charged ion) and an anion (negatively charged ion).

• Electron Transfer: Metals tend to lose valence electrons and become cations; nonmetals tend to gain electrons and become anions.

• Energy Consideration: Ionic bond formation lowers the potential energy of the system.

• Example: Sodium (Na) transfers an electron to chlorine (Cl) to form Na+ and Cl-.

Practice: The strength of an ionic bond comes primarily from the mutual attraction of opposite electrical charges.

Covalent Bonding

Covalent bonding involves the sharing of valence electrons between nonmetal atoms, resulting in the formation of molecules.

• Definition: A bond formed by the sharing of one or more pairs of electrons between atoms.

• Electron Sharing: Atoms achieve stable electron configurations by sharing electrons.

• Example: Two hydrogen atoms share electrons to form H2.

Practice: Elements like argon (Ar) are unlikely to form covalent bonds due to their stable noble gas configuration.

Metallic Bonding

Metallic bonding is characterized by a 'sea' of free-flowing electrons shared among a lattice of metal cations, giving metals their unique properties.

• Definition: The attractive force between free-flowing valence electrons and positively charged metal ions.

• Properties: Responsible for ductility, malleability, luster, and electrical conductivity in metals.

• Example: In copper (Cu), electrons move freely among Cu+ ions.

Practice: Valence electrons that can move freely between metal ions are responsible for metallic bonding.

Electronegativity and Bond Polarity

Electronegativity

Electronegativity (EN) is a measure of an atom's ability to attract electrons in a chemical bond.

• Periodic Trend: Electronegativity increases from left to right across a period and decreases down a group.

• Difference in Electronegativity (ΔEN): Determines the type and polarity of chemical bonds.

• Example: Fluorine (F) is the most electronegative element.

Dipole Moment

A dipole moment arises when there is a significant difference in electronegativity between two bonded atoms, resulting in a polar bond.

• Polarity: Unequal sharing of electrons creates partial charges and a dipole arrow pointing towards the more electronegative atom.

• Example: In CF, the dipole moment points towards F.

Chemical Bond Classifications

The difference in electronegativity between two atoms determines the type of chemical bond:

ΔEN	Bond Type	Bond Illustration
Zero (0)	Pure Covalent	Br–Br
Small (0.1–0.4)	Nonpolar Covalent	C–H
Intermediate (0.5–1.7)	Polar Covalent	Cl–H
Large (>1.7)	Ionic	Na–Cl

Example: S–F bond is more polar than S–Se or S–H.

Nomenclature of Chemical Compounds

Naming Ionic Compounds

Ionic compounds consist of a cation (metal or polyatomic ion) and an anion (nonmetal or polyatomic ion).

• Step 1: The cation is named first and keeps its name.

• Step 2: The anion is named second; if it is a nonmetal, its ending changes to '-ide'.

• Step 3: If a polyatomic ion is present, it keeps its name.

Nonmetal	Base Name
Hydrogen (H)	Hydr-
Nitrogen (N)	Nitr-
Phosphorus (P)	Phosph-
Oxygen (O)	Ox-
Sulfur (S)	Sulf-
Selenium (Se)	Selen-
Tellurium (Te)	Tellur-

Example: CaCl2 is named calcium chloride.

Writing Formulas for Ionic Compounds

• Step 1: Write the ions involved from the compound name.

• Step 2: Use the charges to determine the ratio of ions needed for a neutral compound.

Example: Magnesium sulfate: Mg2+ and SO42– combine to form MgSO4.

Ionic Hydrates

Ionic hydrates are ionic compounds linked to one or more water molecules. The number of water molecules is indicated by a numerical prefix.

Prefix	Number
Mono	1
Di	2
Tri	3
Tetra	4
Penta	5
Hexa	6
Hepta	7
Octa	8
Nona	9
Deca	10

Example: PbO2 · 5 H2O is named lead(IV) oxide pentahydrate.

Naming Molecular (Covalent) Compounds

Molecular compounds consist of nonmetals bonded together. Numerical prefixes indicate the number of each atom present.

• Step 1: The first nonmetal is named normally and uses a numerical prefix if more than one atom is present.

• Step 2: The second nonmetal uses a numerical prefix and its ending changes to '-ide'.

• Note: If the prefix ends with 'a' and the element name starts with a vowel, the 'a' is often dropped.

Prefix	Number
Mono	1
Di	2
Tri	3
Tetra	4
Penta	5
Hexa	6
Hepta	7
Octa	8
Nona	9
Deca	10

Example: Disulfur monochloride: S2Cl

Acids and Acid Nomenclature

Binary Acids

Binary acids consist of hydrogen and one other nonmetal element.

• Step 1: The prefix 'hydro-' is used.

• Step 2: The base name of the nonmetal is used.

• Step 3: The suffix '-ic' is added, followed by 'acid'.

Example: HCl is hydrochloric acid.

Oxyacids

Oxyacids contain hydrogen, oxygen, and another element (usually a nonmetal).

• If the polyatomic ion ends with '-ate', change the ending to '-ic' and add 'acid'.

• If the polyatomic ion ends with '-ite', change the ending to '-ous' and add 'acid'.

Example: H2SO4 (sulfate) becomes sulfuric acid; H2SO3 (sulfite) becomes sulfurous acid.

Key Formulas and Equations

• Formal Charge: $\text{Formal Charge} = (\text{Valence electrons}) - (\text{Nonbonding electrons}) - \frac{1}{2}(\text{Bonding electrons})$

• Difference in Electronegativity: $\Delta EN = |EN_1 - EN_2|$

Summary Table: Bond Types and Properties

Bond Type	Formation	Properties
Ionic	Electron transfer	High melting point, conducts electricity when molten
Covalent	Electron sharing	Low melting point, poor conductor
Metallic	Electron sea	Malleable, ductile, conducts electricity

Additional info: These notes cover topics from General Chemistry chapters on chemical bonding, electronegativity, molecular structure, and nomenclature, suitable for exam preparation.