Ionic Bonds and Ionic Networks

Ionic Bonds, Ions, and Ionic Networks

• Ionic bonds involve electrostatic attraction between oppositely charged ions (metal and non-metal).
• Ions are atoms or molecules gaining/losing electrons, resulting in positive (cations) or negative (anions) charge.
• Ionic network/lattice refers to the 3D arrangement of ions held by ionic bonds.

Formation of Ionic Bonds

• Ionic bond formation starts with a significant electronegativity difference between atoms (metal and non-metal).
• Metals with low ionization energies readily lose valence electrons.
• Non-metals with high electron affinities eagerly accept electrons to attain a full valence shell (octet).
• Metals become positively charged cations, non-metals become negatively charged anions.
• Electrostatic attraction between cation and anion establishes the ionic bond.
• Example: Sodium (Na) donates an electron to Chlorine (Cl), forming Na+ and Cl-.

3D Structure of Ionic Compounds (Ionic Lattice)

• Solid-state ionic compounds form a highly ordered 3D structure called an ionic lattice.
• The lattice has a regular, repeating pattern of positive and negative ions.
• This arrangement maximizes electrostatic attraction and minimizes repulsion.
• Cations are surrounded by anions, and vice versa, forming an alternating charge network.
• This arrangement contributes to a stable, rigid structure with properties like high melting points and brittleness
• The specific arrangement depends on compound stoichiometry and relative ion sizes.

Examples of Ionic Compounds

• NaCl (table salt) has a high melting point and water solubility which makes it valuable for seasoning and industrial processes.
• CaCO3 (limestone/seashells) is hard, making it a useful building material, with basic properties helpful as an antacid.

Covalent Bonds

• Covalent bonds form when atoms share electrons to achieve a stable configuration.
• In single bonds, atoms share one electron pair (e.g., H-H in H2).
• In double bonds, atoms share two electron pairs (e.g., O=O in O2).
• In triple bonds, atoms share three electron pairs (e.g., N≡N in N2).
• The strength of the covalent bond increases with the number of shared electron pairs.

Covalent Molecular Compound Formation

• Non-metal atoms share electrons to attain a full valence shell, forming covalent molecular compounds.
• Shared electrons belong to both atoms, creating a stable electron configuration like a noble gas.
• Hydrogen atoms share electrons with oxygen in water (H2O), allowing all atoms involved to achieve a stable configuration

Physical/Chemical Properties of Covalent Molecular Compounds

• Intramolecular bonds are strong covalent bonds that determine molecule shape/structure.
• Intermolecular forces (IMFs) are weaker forces between molecules, influencing physical properties.
• London Dispersion Forces (LDFs) are the weakest IMFs and become stronger with larger molecules.
• Dipole-Dipole Forces occur between polar molecules and are stronger than LDFs.
• Hydrogen Bonds are the strongest IMFs and involve hydrogen bonded to oxygen, nitrogen, or fluorine.
• Compounds with weak IMFs (gases like oxygen/nitrogen) have low melting/boiling points.
• Stronger IMFs (like water) result in higher melting/boiling points.
• Gases have weak IMFs, liquids have moderate IMFs, and solids have strong IMFs.

Examples of Molecular Compounds

• Hydrogen (H2) is a light, flammable gas used as rocket fuel and in ammonia production due to its high energy content and low density.
• Oxygen (O2) is an essential gas for combustion and is used in welding and medical applications.
• Water (H2O) is a polar solvent with strong hydrogen bonds, dissolving many substances and regulating temperature.
• Carbon dioxide (CO2) , used in carbonated drinks and fire extinguishers, can dissolve under pressure, is unreactive, and can solidify at low temperatures.

Covalent Network Structures

• Covalent bonds involve the sharing of electron pairs between two non-metal atoms.
• Covalent networks (network solids) have atoms interconnected by an extensive network of covalent bonds in a 3D structure.
• Covalent bonds result in a strong attraction holding the atoms together.
• Covalent network solids create one giant molecule.
• Covalent networks exhibit extreme hardness, high melting points, and poor electrical conductivity.

Covalent Bond Formation (Covalent Networks)

• Covalent bonds arise from the mutual attraction of two nonmetal atoms to a shared pair of electrons.
• Valence electron orbitals overlap so electrons can be shared between the nuclei.
• This enables both atoms to reach a stable electron configuration to fill the valence shell.
• Atoms can achieve stability without electron transfer, forming extended network structures.

3D Structure of Covalent Networks

• These compounds have a unique 3D structure characterized by an extended network of covalent bonds connecting all atoms.
• Atoms are interconnected by covalent bonds, forming a large, extended molecule.
• The 3D arrangement is highly ordered and rigid which create a strong, inflexible structure.

Electrons in Covalent Network Structures

• Electrons are shared between atoms forming a region of high electron density that attracts positively charged nuclei.
• Attractions result in strong covalent bonds that hold atoms together.
• The sharing of electrons allows all atoms to achieve a stable electron configuration.

Properties of Covalent Networks

• Simple molecules have weak intermolecular forces, resulting in low melting/boiling points.
• Covalent networks are interconnected by covalent and have very different electrical properties.
• They exhibit a rigid, 3D structure contributing to extreme hardness and high melting points.
• Absence of free electrons in the network leads to poor electrical conductivity.

Examples of Covalent Network Compounds

• Diamond (C): Each carbon is covalently bonded to four others in a tetrahedral arrangement which gives diamond its hardness (used for cutting tools).
• Silicon dioxide (SiO2) i.e quartz, has abundance and hardness making it useful in construction and glass.

Metallic Bonding

• Metallic bonding is unique since it involves the delocalization of valence electrons, creating a "sea of electrons."
• Metal atoms lose valence electrons and arrange as positively charged metal ions in a lattice.
• Valence electrons move freely throughout the lattice, not bound to any specific atom.
• Electrostatic attraction between metal ions and electrons constitutes the metallic bond that is non-directional.
• The "sea of electrons" contributes to shininess since electrons absorb/re-emit light.

Formation of Metallic Bonds

• Metal atoms lose valence electrons and don't become attached but move freely.
• This forms a "sea of electrons" around positively charged electrons.
• Attraction between the "sea" and positive ions creates the nature of the bond.
• The non-directional bond enables atoms to slide past each other, giving malleability and ductility.

Role of Electrons/Ions in Metallic Bonds

• Electrons/ions are crucial in the formation and properties of metals.
• Metals lose electrons to form positive ions and allows them to float in the "sea".
• Electrons are responsible for high electrical/thermal conductivity since the bond doesn't break.

Metallic Compound Examples

• Iron (Fe) is used in structural applications due to its strength and malleability.
• Copper (Cu) is used in electrical wiring due to its conductivity and malleability.
• Aluminum (Al) is used in aircraft construction due to its low density and high strength.
• Gold (Au) is used in jewelry due to its conductivity and corrosion resistance.

Hydrogen Bonding

• It is a strong dipole-dipole attraction between a hydrogen atom and a highly electronegative atom (F, O, N).
• Electronegativity polarizes a molecule since electrons are pulled towards electronegative atom which gives the hydrogen atom a slight positive charge.
• The Hydrogen bonds are just between molecules, not within.

Hydrogen Bond Formation

• Created through an interaction between positive hydrogen atoms and a negative atom.
• Electronegativity difference polarizes the bond.
• A dipole moment in the molecule is created since partial charges create ends.
• Partially positive hydrogen bonds with a lone pair on another electronegative atom.
• Line-up will maximize electrostatic attraction.

Role of Electrons in Hydrogen Bonding

• Electrons pull electron density away from the hydrogen atom (becomes positive).
• Molecules attract, the charges creates a dipole moment with negative electrons drawing in the atom.

Properties of Hydrogen Bonding

• Compounds exhibiting hydrogen bonding have unique physical properties due to strong intermolecular forces.
• Hydrogen bonds increase boiling/melting points since energy is needed to overcome these connections.
• Compounds with bonds will be soluble in polar solvents.

Examples of Hydrogen Bonding

• Water (H2O) has hydrogen atoms bonded to oxygen, giving it a high boiling point, surface tension, and solvent properties.
• Alcohols (like ethanol) have hydrogen bonds, allowing a solvent in various applications.