Ionic bonds occur when one atom transfers electrons to another, resulting in the formation of charged ions. This electron transfer is typical between metals and nonmetals.

Covalent bonds form when atoms share electron pairs, allowing each atom to attain a stable electron configuration. This sharing mechanism is fundamental to the structure of molecules.

Ionic compounds conduct electricity when molten because the ions are free to move. This distinguishes them from many covalent compounds which do not conduct electricity in their solid forms.

Metallic bonding involves electrons that are not bound to any one atom but move freely throughout the structure. This 'sea of electrons' is what allows metals to conduct electricity efficiently.

Ionic compounds have a rigid lattice structure that can fracture when subjected to stress, making them brittle. In contrast, metals typically exhibit malleability due to their bonding structure.

Sodium, a metal with a low ionization energy, readily loses an electron, while chlorine, a nonmetal with high electron affinity, readily gains an electron. This difference in properties leads to the formation of an ionic bond.

The polarity of a covalent bond is determined by the difference in electronegativity between the bonded atoms. A larger difference leads to an uneven distribution of electron density, making the bond polar.

A molecule with four bonding pairs and no lone pairs adopts a tetrahedral geometry to minimize electron pair repulsions. This geometry is one of the simplest predictions of the VSEPR theory.

Triple bonds consist of one sigma bond and two pi bonds, which together require more energy to break than single or double bonds. This higher bond energy means that triple bonds are among the strongest types of covalent bonds.

A coordinate covalent bond is formed when one atom donates both electrons needed for the bond, unlike typical covalent bonds where each atom contributes one electron. This donation leads to a bond that is structurally similar to other covalent bonds.

A greater difference in electronegativity causes electrons to be more unequally shared, which increases the ionic character of the bond. This shift in electron distribution is key to understanding bond properties.

Simple covalent compounds usually have low melting and boiling points due to weak intermolecular forces. High melting and boiling points are more characteristic of network covalent solids rather than simple molecular compounds.

Water has two lone pairs in addition to its bonding pairs around the oxygen atom, causing a bent shape. This geometry minimizes repulsion between all electron pairs in accordance with VSEPR theory.

Covalent bonds involve the sharing of electrons in specific orbitals, which makes them directional and key in determining the shape of molecules. Other bond types, such as ionic and metallic bonds, do not impose strict geometric directions.

Bond energy is defined as the energy required to break one mole of bonds in a gaseous substance. It serves as a measure of the strength and stability of a chemical bond.

Ionic bond formation is most favorable when one atom (typically a metal) has a low ionization energy, enabling it to lose electrons easily, and the other atom (typically a nonmetal) has a high electron affinity, favoring electron gain. This combination leads to the formation of oppositely charged ions that attract each other.

A double bond is made up of one sigma bond and one pi bond, which increases the electron density between the two nuclei. This additional electron sharing makes a double bond stronger than a single bond, which only has one sigma bond.

Resonance involves the delocalization of electrons across multiple bonds, which effectively averages the bond order. This results in bond lengths that are neither as long as those of a single bond nor as short as double bonds, enhancing molecular stability.

Molecular orbital theory shows that O2 has a bond order of 2, resulting from the combination of sigma and pi bonding interactions. Additionally, the presence of unpaired electrons in its molecular orbitals makes O2 paramagnetic.

A dative or coordinate covalent bond is formed when one atom supplies both electrons for the bonding pair, which distinguishes it from a traditional covalent bond where each atom contributes one electron. This subtle difference is important in the formation of complex molecules and coordination compounds.