Chemistry Unit 1.4

Vocabulary flashcards reviewing types of chemical bonding, electronegativity, intermolecular forces, and VSEPR molecular shapes with corresponding electron pair counts and bond angles.

Covalent Bond

A bond formed between two non-metals where electrons are shared between outer electron shells to achieve a full outer shell, with balanced electrostatic forces.

Coordinate (Dative) Bond

A bond formed when both electrons in the shared pair are supplied from a lone electron pair of a single atom; once formed, it acts and reacts in exactly the same way as a standard covalent bond.

Electronegativity

The power of an atom to attract negative charge towards itself within a covalent bond.

Periodic Trends in Electronegativity

Electronegativity increases across a period because the atomic radius decreases, and it decreases down a group because electron shielding increases.

Intermediate Bond

A bond in which electrons are shared unequally due to a difference in electronegativity, creating partial charges ($\delta^+$ and $\delta^-$) on the bonded atoms.

Ionic Bond

A bond characterized by a complete transfer of electrons between atoms due to a very large electronegativity difference, producing distinct ions with full charges.

Permanent Dipole

A lasting separation of electric charge within a polar covalent bond where the more electronegative atom pulls negative charge toward itself, creating a $\delta^-$ region and a $\delta^+$ region.

Induced Dipole

A temporary dipole created when the electron orbitals around a molecule are distorted or influenced by an approaching charged particle.

Van der Waals Forces

The weakest type of intermolecular force, acting between molecules and varying in strength depending on the relative molecular mass ($M_r$) and molecular shape.

Effect of Molecular Shape on Van der Waals Forces

Straight-chain molecules experience stronger van der Waals forces than branched-chain molecules because they can pack closer together, reducing the distance over which the forces act.

Hydrogen Bonding

The strongest type of intermolecular force, occurring when hydrogen is bonded to nitrogen, oxygen, or fluorine, and interacts with a lone electron pair on one of these electronegative atoms in another molecule.

VSEPR Principle

Valence Shell Electron Pair Repulsion principle; states that electron pairs around a central atom repel each other to adopt positions that produce the largest possible bond angle, with lone pairs providing additional repulsion that reduces bond angles between covalent bonds by $2.5^\circ$ per lone pair.

Electron Pair Repulsion Hierarchy

The relative strength of repulsion between electron pairs: $\text{lone pair - lone pair} > \text{lone pair - bonding pair} > \text{bonding pair - bonding pair}$.

Linear Molecular Shape

A molecular geometry with $2$ bonding pairs and $0$ lone pairs around the central atom, resulting in a bond angle of $180^\circ$ (e.g., $BeCl_2$).

Trigonal Planar Molecular Shape

A molecular geometry with $3$ bonding pairs and $0$ lone pairs around the central atom, resulting in a bond angle of $120^\circ$ (e.g., $BF_3$, $BCl_3$).

Tetrahedral Molecular Shape

A molecular geometry with $4$ bonding pairs and $0$ lone pairs around the central atom, resulting in a bond angle of $109.5^\circ$ (e.g., $CH_4$, $SiCl_4$).

Pyramidal Molecular Shape

A molecular geometry with $3$ bonding pairs and $1$ lone pair around the central atom, resulting in a bond angle of $107^\circ$ (e.g., $NH_3$, $PCl_3$).

V-Shaped Molecular Shape

A bent molecular geometry with $2$ bonding pairs and $2$ lone pairs around the central atom, resulting in a bond angle of $104.5^\circ$ (e.g., $H_2O$).

Trigonal Bipyramidal Molecular Shape

A molecular geometry with $5$ bonding pairs and $0$ lone pairs around the central atom, resulting in bond angles of $120^\circ$ and $90^\circ$ (e.g., $PF_5$, $PCl_5$).

Octahedral Molecular Shape

A molecular geometry with $6$ bonding pairs and $0$ lone pairs around the central atom, resulting in bond angles of $90^\circ$ (e.g., $SF_6$).