Bonding: General Concepts – Structured Study Notes for General Chemistry

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Bonding: General Concepts

Types of Chemical Bonds

Chemical bonds are the forces that hold atoms together in compounds. The three primary types of chemical bonds are covalent, ionic, and metallic bonds.

Covalent Bonds: Occur when electrons are shared between two or more nonmetals. The attraction between the electrons of one atom and the nucleus of another leads to bond formation if the attraction outweighs repulsion. Resulting compounds are called molecules or molecular compounds.
Ionic Bonds: Formed when electrons are transferred from a metal (which loses electrons to become a cation) to a nonmetal (which gains electrons to become an anion). Ionic compounds are typically formed between metals and nonmetals.
Metallic Bonds: Involve the sharing of electrons among a lattice of metal atoms (not covered in detail in this chapter).

Representation of Molecular Structures: Molecules can be represented by chemical formulas, structural formulas, space-filling models, and ball-and-stick models.

3D molecular model Attractive and repulsive forces in bond formation Methane molecule representations: structural, space-filling, ball-and-stick

Bond Length

Bond length is the distance between two bonded atoms where the energy of the system is lowest. This optimal distance results from a balance between attractive and repulsive forces.

Potential energy curve for H-H bond formation

Ionic Bonds: Formation and Properties

Ionic bonds occur when an atom with low ionization energy (metal) transfers electrons to an atom with high electron affinity (nonmetal). The resulting ions are held together by electrostatic attraction.

Cation: A positively charged ion formed by loss of electrons (e.g., Na+).
Anion: A negatively charged ion formed by gain of electrons (e.g., Cl-).

Formation of sodium ion from sodium atom

Categories of Elements

Metals: Shiny, solid, good conductors, malleable, ductile. Exception: Mercury is liquid.
Nonmetals: Can be solids (brittle), liquids, or gases; nonconductors.
Metalloids: Shiny like metals but brittle like nonmetals; semiconductors.

Periodic table of elements

Commonly Used Names of Groups

Groups in the periodic table are often referred to by specific names:

Group 1: Alkali metals
Group 2: Alkaline earth metals
Group 16: Chalcogens
Group 17: Halogens
Group 18: Noble gases

Periodic table highlighting common groups

Electronegativity and Bond Polarity

Electronegativity

Electronegativity (χ) is a measure of an atom's ability to attract electrons in a chemical bond. The difference in electronegativity between two atoms determines the bond's polarity:

Δχ < 0.4: Nonpolar covalent bond
0.4 < Δχ < 2.0: Polar covalent bond
Δχ > 2.0: Ionic bond

Periodic table with electronegativity values

Unequal Sharing and Polar Bonds

In covalent bonds, electrons may be shared unequally, resulting in a polar covalent bond with partial positive and negative charges (dipole).

Charge distribution in covalent and ionic bonds

Nonpolar Covalent Bonds

Nonpolar covalent bonds occur when electrons are shared equally, typically between identical atoms or when the electronegativity difference is very small (Δχ ≤ 0.4).

Nonpolar covalent bond: Cl2

Polar Covalent Bonds

Polar covalent bonds occur when electrons are shared unequally due to a moderate difference in electronegativity (0.4 < Δχ < 2.0). This results in a dipole moment.

Polar covalent bond: HCl Dipole moment in HCl

Ions: Electron Configurations and Sizes

Electron Configuration of Compounds

Atoms achieve stable electron configurations (often noble gas configurations) through electron sharing (covalent) or transfer (ionic). Isoelectronic species have the same number of electrons.

Predicting Formulas of Ionic Compounds

Ionic compounds are electrically neutral. The empirical formula is determined by balancing the charges of the ions involved.

Sizes of Ions

Ion size affects the structure and stability of ionic solids. Isoelectronic ions are compared based on nuclear charge; the ion with the smallest nuclear charge is largest.

The Covalent Chemical Bond: A Model

Bonding Models

Models are used to explain chemical bonding at the microscopic level. They are simplifications and may evolve over time as new data emerges.

The Localized Electron Bonding Model

LE Model

The Localized Electron (LE) Model describes molecules as collections of atoms bound by shared pairs of electrons. Lone pairs are localized on atoms, while bonding pairs are shared between atoms.

Lewis Structures

Writing Lewis Structures

Lewis structures show the arrangement of valence electrons in molecules. Only valence electrons are included, and dots represent electrons. The octet rule states that atoms form stable molecules when surrounded by eight electrons.

Draw the basic molecule with single bonds.
Sum the valence electrons.
Satisfy the octet for outer atoms.
Place leftover electrons on the central atom.
If the central atom is not satisfied, form double bonds.
For charged molecules, use brackets and indicate the charge.

Example: Lewis Structures for Common Molecules

Table of Lewis structures for HF, N2, NH3, CH4, CF4, NO+

Exceptions to the Octet Rule

Incomplete Octet

Boron and beryllium often form compounds with fewer than eight electrons around them.

Expanded Octet

Elements in Period 3 and beyond can have more than eight electrons (expanded octet), as seen in molecules like SF6 and PCl5.

Lewis structure for SF6

Resonance

Resonance Structures

Some molecules have more than one valid Lewis structure. The true structure is a hybrid of all resonance forms, represented by double-headed arrows.

Resonance structures for nitrate ion Resonance hybrid for nitrate ion

Formal Charge

Calculating Formal Charge

Formal charge helps estimate the charge distribution in molecules. It is calculated as:

Formal charge = (number of valence electrons on free atom) − (number of electrons assigned in molecule)
Atoms strive for formal charges close to zero; negative charges should reside on the most electronegative atoms.

Naming Simple Compounds (Nomenclature)

Binary Ionic Compounds (Type I)

Type I compounds have a cation followed by an anion. The cation is named first, and the anion is named by adding -ide to the root of the element name.

Common Monatomic Cations and Anions

Name	Anion	Name	Cation
Hydride	H−	Hydrogen	H+
Fluoride	F−	Lithium	Li+
Chloride	Cl−	Sodium	Na+
Bromide	Br−	Potassium	K+
Iodide	I−	Cesium	Cs+
Oxide	O2−	Beryllium	Be2+
Sulfide	S2−	Magnesium	Mg2+
Nitride	N3−	Calcium	Ca2+
Phosphide	P3−	Barium	Ba2+
		Aluminum	Al3+