Chemical Bonding: Molecular & Ionic Compounds and Lewis Structures

Chemical Bonding: Molecular & Ionic Compounds

Introduction to Chemical Bonding

Chemical bonding occurs because atoms seek to achieve the lowest possible energy state. A chemical bond forms if the system can lower its total energy, primarily through the net potential energy resulting from the attraction between charged particles and the kinetic energy of their motion.

• Ionic Bonds: Formed between metals and nonmetals; electrons are transferred from the metal to the nonmetal.

• Covalent Bonds: Formed between nonmetals; electrons are shared between atoms.

• Metallic Bonds: Formed between metals; electrons are pooled and delocalized over a lattice of metal atoms.

Types of Chemical Bonds

The type of bond formed depends on the elements involved and their electronegativity differences.

Atoms	Type of Bond	Characteristics of Bond
Metal + Nonmetal	Ionic	Electrons transferred
Nonmetal + Nonmetal	Covalent	Electrons shared
Metal + Metal	Metallic	Electrons pooled

Bond Formation and Energy

• Ionic Bonding: Metals lose electrons (form cations), nonmetals gain electrons (form anions). Oppositely charged ions attract, lowering potential energy (Coulomb's Law).

• Covalent Bonding: Nonmetals have high ionization energy and tend to share electrons, which interact with both nuclei, lowering potential energy.

• Metallic Bonding: Metals have low ionization energy and lose electrons easily. Valence electrons are pooled and delocalized over the metal lattice.

Representing Valence Electrons: Lewis Symbols

Lewis Dot Symbols

The Lewis symbol represents the valence electrons of main-group elements as dots around the element's symbol. Each dot stands for a valence electron, with a maximum of 8 dots (2 per side).

• Example: Oxygen: $1s^2 2s^2 2p^4$ → 6 valence electrons, shown as 6 dots around the symbol "O".

• Octet Rule: Atoms tend to achieve a stable configuration with eight electrons in their valence shell.

Electron Transfer and Lewis Structures

In ionic bonding, electrons are transferred from the metal to the nonmetal. Lewis symbols can be used to illustrate this process and predict chemical formulas.

• Example: Potassium and Chlorine: $\mathrm{K} + \mathrm{Cl} \rightarrow \mathrm{K}^+ + \mathrm{Cl}^-$ Lewis symbols: $:\mathrm{K}: + :\mathrm{Cl}: \rightarrow \mathrm{K}^+ [:\mathrm{Cl}:]$

• Example: Sodium and Sulfur: $2\mathrm{Na} + \mathrm{S} \rightarrow \mathrm{Na}_2\mathrm{S}$ Lewis symbols: $2:\mathrm{Na}: + :\mathrm{S}: \rightarrow 2\mathrm{Na}^+ [:\mathrm{S}:]^{2-}$

Electron Configuration and Ion Formation

• Example: Sodium ion: $[\mathrm{Ne}] 3s^1 \rightarrow [\mathrm{Ne}]$

• Example: Sulfide ion: $[\mathrm{Ne}] 3s^2 3p^4 \rightarrow [\mathrm{Ne}] 3s^2 3p^6$ or $[\mathrm{Ar}]$

Using Lewis Symbols to Predict Formulas

• Example: Calcium and Chlorine: $\mathrm{Ca}^{2+} + 2[:\mathrm{Cl}:]^- \rightarrow \mathrm{CaCl}_2$

• Example: Magnesium and Nitrogen: $3\mathrm{Mg}^{2+} + 2[:\mathrm{N}:]^{3-} \rightarrow \mathrm{Mg}_3\mathrm{N}_2$

Key Terms and Concepts

• Valence Electrons: Electrons in the outermost shell, involved in bonding.

• Octet Rule: Atoms tend to have eight electrons in their valence shell for stability.

• Cation: Positively charged ion (metal loses electrons).

• Anion: Negatively charged ion (nonmetal gains electrons).

• Lewis Structure: Diagram showing the arrangement of valence electrons among atoms in a molecule or ion.

Examples and Applications

• Table Salt (NaCl): Ionic compound formed by transfer of electrons from Na to Cl.

• Water (H2O): Covalent compound formed by sharing electrons between H and O.

• Sodium Metal (Na): Metallic bonding with delocalized electrons.

Additional info: The notes provide foundational concepts for understanding chemical bonding, including the use of Lewis structures to represent electron arrangements and predict compound formulas. These principles are essential for further study in General Chemistry, especially in topics such as molecular geometry, polarity, and reactivity.