Chemical Bonding Practice Flashcards

A set of vocabulary flashcards covering chemical bonding concepts, including bond types, molecular geometry, and intermolecular forces based on the lecture transcript.

Ionic bond

The electrostatic attraction between oppositely charged ions (cations and anions).

Metallic bond

The electrostatic attraction between cations and delocalized electrons.

Covalent bond

The electrostatic attraction between the nuclei of two atoms and a shared pair of electrons.

Co-ordinate/dative covalent bond

A type of covalent bond where the shared pair of electrons comes from the same atom.

Electronegativity

The power of an atom that is covalently bonded to another atom to attract the bonding pair of electrons to itself.

Bond energy

The energy required to break one mole of a specific bond in a molecule that is in the gaseous state.

Metallic bonding strength

The strength increases with increasing positive charge on ions, decreasing size of metal ions, and an increasing number of mobile electrons per atom.

Bonding pair

A shared pair of electrons that are involved in covalent bonding.

Lone pair

Pairs of electrons in the outer shell of an atom that are not involved in bonding.

Incomplete octet

A state where an atom in a molecule has fewer than eight electrons in its valence shell, such as in $BF_3$ or $BeCl_2$.

Expanded octet

A state where an atom in a molecule has more than eight electrons in its valence shell, such as in $SF_6$, $PF_5$, or $XeF_4$.

Molecular orbital

A combined orbital formed by two atomic orbitals overlapping to form a covalent bond.

Pauli exclusion principle

The requirement that the maximum number of electrons in an orbital is two, and they must have opposite spins.

Sigma ($\sigma$) bond

A covalent bond formed by the end-to-end overlapping of atomic orbitals from two different atoms.

Pi ($\pi$) bond

A covalent bond formed by the side-by-side overlapping of two $p$ orbitals; these are weaker than sigma bonds.

Orbital hybridisation

The process where atomic orbitals participate in covalent bonds after being singly occupied and mixing into new hybrid states like $sp$, $sp^2$, or $sp^3$.

VSEPR Theory

Valence Shell Electron Pair Repulsion Theory, used to deduce molecular shapes based on the least repulsion between bonding and lone pairs of electrons.

Bond length

The internuclear distance between two atoms in a covalent bond.

Bond polarity

The partial separation of charge resulting from the unequal sharing of a bonding pair of electrons between two different atoms.

Dipole moment

The vector sum of the dipole moments of all the individual bonds in a molecule.

Instantaneous dipole-induced dipole forces

Attractive intermolecular forces caused by temporary dipoles; they are present in all molecules and their strength depends on the number of electrons and molecular size.

Permanent dipole-permanent dipole forces

Attractive intermolecular forces caused by constant dipoles in polar molecules.

Hydrogen bonding

The strongest intermolecular force, occurring when hydrogen is covalently bonded to $F$, $O$, or $N$ and is attracted to the lone pair of $F$, $O$, or $N$ on a second molecule.

Surface tension

Attractive forces of molecules at the surface of a liquid that allow it to resist an external force and reduce its surface area.