Intramolecular Forces

Ionic Bonds

Definition: Ions of opposite charges held together by electrostatic attraction. 

Physical State: Forms crystal lattice structures of ions. 

Polarity: Ionic bonds are very polar. 

Solubility: Has relatively high solubility in polar solvents, where it can dissociate into ions. 

Intermolecular Forces: Strong intermolecular forces due to electrostatic attraction between ions of opposite charges. 

Conductivity: 

• As a liquid or in an aqueous state: Conductive due to freely-moving ions from overcoming intermolecular forces. 

• As a solid: Not conductive, as ions are locked in place. 

Physical Property: Hard yet brittle, as shifting the ions' positions easily results in repulsion. 

Covalent Bonds

Definition: Chemical bonds that involve the sharing of valence electrons. 

• Each bond consists of two electrons. 

Types of Covalent Bonds: 

• All covalent bonds have a σ-bond, which is a direct bond between two orbitals facing each other on the same axis. 

  • Single bonds: Consists of one σ-bond. Total of two valence electrons shared. 

• π-bonds are formed by the side-by-side overlap of p-orbitals, which are perpendicular to the axis connecting the nuclei of the bonded atoms. 

  • Double bonds: Consists of one σ-bond and one π-bond. Total of four valence electrons shared. 

  • Triple bonds: Consists of one σ-bond and two π-bond. Total of six valence electrons shared. 

Polarity: Covalent bonds can be either polar or nonpolar, depending on the electronegativity difference between the bonded atoms of compounds. 

Intermolecular Forces: Generally has weaker intermolecular forces than most ionic and metallic compounds. 

Conductivity: Usually poor conductors, as most covalent compounds don't dissociate or possess delocalized electrons. 

Physical States: 

• Molecular Compounds: Usually exists as a gas, a liquid, or a solid with a low melting-point. 

• Covalent Network Solids: Usually exist as solids with high melting points. Unlike molecular compounds, which are composed of discrete units, they form a continuous network of covalently bonded atoms. 

Metallic Bonds

Definition: Occur between metal atoms where electrons are free to move throughout a lattice of positive ions. 

Intermolecular Forces: Strong intermolecular forces due to electrostatic attraction between the ions and the 

Polarity: Metallic bonds are considered nonpolar due to the delocalized, nondirectional "sea" of electrons. 

Solubility: Depends on reactivity. More reactive metals may undergo redox reactions with solvents, forming aqueous ions. 

Conductivity: Highly conductive due to the delocalized electrons within the metal. 

Physical Property: 

• Malleable and ductile due to its delocalized electrons. 

• Has relatively higher luster. More inert metals reflect more light due to resistance from oxidation. 

Metallic Alloys: 

• Substitutional Alloy: Metallic elements with similar atomic radii replace atoms in the host metal's lattice, forming a relatively uniform structure. (e.g. brass, composed of copper and zinc)

• Interstitial Alloy: Smaller atoms fit into the spaces between larger metal atoms in the lattice. (e.g. steel, consisting of iron and carbon)