Bonding_and_structure_pmt

Topic 2: Bonding and Structure

Ionic Bonding

• Ionic bonding is a type of chemical bond that occurs between a metal and a non-metal. Metals tend to lose electrons, while non-metals tend to gain electrons, leading to the transfer of electrons from the metal to the non-metal.

• This electron transfer results in the formation of charged particles known as ions, wherein the metal becomes positively charged (cation) and the non-metal becomes negatively charged (anion).

• A common example of ionic bonding is sodium chloride (NaCl). In this case, sodium (Na) loses one electron to become Na⁺, while chlorine (Cl) gains an electron to become Cl⁻.

• The oppositely charged ions attract each other through electrostatic forces, forming a crystal lattice structure known as a giant ionic lattice. This arrangement is characterized by its regular, repeating arrangement of ions.

• The strength of the ionic bond is influenced by several factors:

  • Higher ion charge results in stronger forces of attraction, leading to a stronger ionic bond.

  • The ionic radius plays a role; larger ions have a weaker attraction due to increased distance from the charged nucleus.

• Dot and cross diagrams are tools used to visualize ionic bonding, where the transferred electrons are represented by dots or crosses to distinguish which atom they originated from.

Covalent Bonding

• Covalent bonds occur between two non-metals and involve the sharing of electrons to achieve stable electron configurations, often filling the outer electron shells of the bonding atoms.

• In a covalent bond, there exists a strong electrostatic attraction between the positively charged nuclei and the shared electrons.

• Covalent bonding can involve single, double, or triple bonds, depending on the number of shared electron pairs:

  • A single bond represents one pair of shared electrons (denoted as -).

  • A double bond involves two pairs of shared electrons (denoted as =).

  • A triple bond consists of three pairs of shared electrons (denoted as ≡).

• Dot and cross diagrams can also represent covalent bonding, illustrating how electron pairs are shared between atoms.

• The strength of covalent bonds is often inversely related to bond length; shorter bonds typically exhibit greater strength due to increased overlap of atomic orbitals.

Dative Bonding (Coordinate Bonding)

• Dative bonds, also known as coordinate bonds, occur when one atom donates both electrons in a shared pair. This type of bonding is particularly important in complex ions and certain molecular entities.

• A prime example is the formation of the ammonium ion (NH₄⁺), where a hydrogen ion (H⁺) accepts a lone pair from an ammonia molecule (NH₃), resulting in a stable cation.

• Once formed, dative bonds behave similarly to standard covalent bonds and are represented in diagrams using standard representations of electron sharing (either both dots or both crosses).

Properties of Simple Covalent Substances

• Simple molecular structures consist of discrete covalently bonded molecules that are held together by relatively weak intermolecular forces known as van der Waals forces.

• Common examples include water (H₂O) and iodine (I₂).

• Due to these weak intermolecular forces, substances with simple molecular structures typically exhibit:

  • Low melting and boiling points, as less energy is required to overcome the intermolecular forces during phase changes.

  • Poor electrical conductivity, as they lack free-moving charged particles that contribute to electrical conduction.

Shapes of Simple Molecules

• The three-dimensional geometry of simple molecules is determined by the number of electron pairs (bonding and lone pairs) surrounding the central atom.

• Lone pairs exert greater repulsion than bonding pairs, causing alterations in bond angles. Common molecular shapes and their associated bond angles include:

  • Linear: 2 bonding pairs, 0 lone pairs → 180° angle.

  • Trigonal Planar: 3 bonding pairs, 0 lone pairs → 120° angle.

  • Trigonal Pyramid: 3 bonding pairs, 1 lone pair → 107° angle.

  • Tetrahedral: 4 bonding pairs, 0 lone pairs → 109.5° angle.

  • Trigonal Bipyramidal: 5 bonding pairs → 90° and 120° angles.

  • Octahedral: 6 bonding pairs → 90° angle.

Bond Polarity and Electronegativity

• Electronegativity is defined as the ability of an atom to attract electrons in a bond. It generally increases across a period in the periodic table and decreases down a group due to changes in atomic size and shielding effects.

• Bond polarity arises from differences in electronegativity between bonded atoms, leading to the formation of polar and nonpolar bonds:

  • Polar bonds occur when there is a significant difference in electronegativity, resulting in dipoles (partial positive and negative charges).

  • Polar molecules exhibit overall polarity, which affects their interactions with other substances. An illustrative example is water (H₂O), which is polar due to its bent molecular shape and the presence of polar covalent bonds.

Intermolecular Forces

• Van der Waals Forces: Represent the weakest type of intermolecular forces that arise from temporary induced dipoles when molecules are close to each other. The strength of these forces varies with molecular size and shape; for instance, longer molecular chains exhibit stronger forces.

• Permanent Dipoles: These occur in molecules with polar bonds, where the differences in electronegativity create regions of partial positive and negative charges, leading to strong intermolecular attractions.

• Hydrogen Bonding: The strongest type of intermolecular force, which occurs specifically when hydrogen is bonded to highly electronegative atoms (nitrogen, oxygen, or fluorine). The hydrogen atoms form strong attractions with the electron-rich regions of the electronegative atoms, resulting in significantly elevated melting and boiling points for compounds such as water.

Properties of Bonding Types

• Ionic Compounds: Known for their high melting points and ability to conduct electricity when melted or dissolved in water due to the presence of free-moving ions.

• Metallic Compounds: Excellent electrical conductors due to the presence of delocalized electrons, which can move freely throughout the structure. They also possess malleability and ductility due to the non-directional nature of metallic bonding.

• Simple Molecular Compounds: Characterized by low melting points and poor electrical conductivity, as these compounds lack mobile charged particles.

• Macromolecular Structures: Exhibit very high melting points due to the presence of strong covalent bonds that require significant energy to break. Notable examples include diamond and graphite.

  • Diamond: An extremely hard material where each carbon atom is tetrahedrally bonded to four other carbon atoms, forming a rigid three-dimensional structure.

  • Graphite: Consists of layers of carbon atoms arranged in a hexagonal lattice; each layer allows free-moving delocalized electrons, which impart electrical conductivity.

Graphene

• Graphene is defined as a single layer of graphite, notable for its exceptional strength and lightweight properties. It also conducts electricity very effectively due to the presence of delocalized electrons, making it a significant material in various technological applications and research fields.