Bonding and Structure - Vocabulary Flashcards

Bonding

  • Bonding is the chemical combination of atoms or elements to form compounds. The force of attraction holding atoms or elements together in a molecule/crystal is called a chemical bond.

  • Bonding/combination occurs mainly in four forms:

  • Ionic/electrovalent bonding – transfer of electrons from a metal atom to a non-metal atom; occurs between metals and non-metals.

    1. Covalent bonding – sharing of electrons between two or more non-metal atoms/elements; bonding electrons are contributed by the participating atoms.

    2. Dative/co-ordinate bonding – sharing of bonding electrons donated by one atom/molecule involved.

    3. Metallic bonding – attraction between metal ions and the valence electrons within the metal lattice.

Ionic/Electrovalent Bonding

  • Involves transfer of electrons from a metal to a non-metal.

  • The number of electrons lost by the metal equals its valency; the number of electrons gained by the non-metal equals its valency.

  • Formation of ions:

    • Metal atom loses electrons to form a positively charged ion (cation).

    • Non-metal atom gains electrons to form a negatively charged ion (anion).

  • The oppositely charged ions attract each other; the force of attraction is the electrovalent/ionic bond.

  • Compounds formed are ionic/electrovalent compounds.

  • Metals lose electrons to achieve noble gas electronic configuration; non-metals gain electrons to achieve noble gas configuration.

  • Example: Formation of sodium chloride

    • When sodium reacts with chlorine: ext{Na}
      ightarrow ext{Na}^+ + e^-

    • Chlorine gains an electron: ext{Cl} + e^-
      ightarrow ext{Cl}^-

    • Result: ext{Na}^+ ext{ and } ext{Cl}^- form an ionic bond to give ext{NaCl}

    • Diagrammatic illustration (conceptual): Na (2:8:1) and Cl (2:8:7) combine to form a stable ionic lattice.

  • Example: Calcium chloride

    • During formation, calcium loses two electrons to form ext{Ca}^{2+}; two chlorine atoms gain one electron each to form two ext{Cl}^- ions.

    • Resultant species: ext{Ca}^{2+} ext{ and } 2 ext{Cl}^-, held together by ionic bonds.

    • Formulation: ext{CaCl}_2

Covalent Bonding

  • Occurs between non-metal elements/atoms to form compounds.

  • Involves mutual sharing of electrons; each atom contributes electrons so that sharing results in a stable electronic configuration like noble gases.

  • The compounds formed are covalent compounds.

  • Covalent bonds can be classified as single, double, triple, or quadruple depending on the number of electrons shared.

  • a) Single covalent bond

    • Formed when one pair of electrons is shared (each atom contributes one electron).

    • Examples:

    • i) Hydrogen molecule: ext{H}_2 ext{ (H–H)}

    • ii) Water molecule: ext{H}_2 ext{O} with H–O–H; each H shares one electron with O.

    • iii) Ammonia: ext{NH}_3 (N with three H atoms)

    • iv) Methane: ext{CH}_4 (C bonded to four H atoms)

    • Note: The outer energy level lone pair (non-bonding pair) on atoms not involved in bonding is called a lone pair.

  • b) Double covalent bond

    • Formed when two pairs of electrons are shared between two bonded atoms.

    • Example:

    • i) Oxygen molecule: ext{O}_2 (O=O)

    • ii) Carbon dioxide: ext{CO}_2 (O=C=O)

  • c) Other covalent considerations

    • Covalent bonding often leads to the formation of simple molecular structures (molecules as discrete units) or giant molecular structures (extended networks).

  • Dative/Coordinate bonding

    • A type of covalent bond where both electrons in the bond come from the same atom.

    • The donor atom provides a lone pair to form the bond with the acceptor.

Metallic Bonding

  • Occurs between atoms of metal elements.

  • In a metal lattice, valence electrons are released into a general pool (delocalized electrons) while metal ions remain fixed in the lattice.

  • The electrostatic attraction between the positively charged metal ions and the delocalized electrons holds the structure together.

  • Strength of metallic bonds increases with the number of electrons released into the electron cloud.

  • Example: Metals like iron and aluminum release up to three electrons each to the electron pool, resulting in a strong metallic bond; metals like sodium and potassium release only one electron each, resulting in comparatively weaker bonding.

Structure of Compounds/Substances

  • Types of bonding give rise to different structural types:

    • Ionic/electrovalent → Giant ionic structure

    • Covalent → Simple molecular structure or Giant molecular structure

    • Metallic → Giant metallic structure

  • A giant ionic structure is a three-dimensional crystal lattice of alternating cations and anions.

  • In such lattices, each ion is surrounded by several oppositely charged ions; the number of surrounding ions is the coordination number.

  • Example structures include: ext{NaCl} and ext{MgCl}_2 with a canonical 6:6 coordination in NaCl.

  • Example: Structure of sodium chloride

    • Each ext{Na}^+ ion is surrounded by six ext{Cl}^- ions and each ext{Cl}^- is surrounded by six ext{Na}^+ ions → coordination number 6:6.

Properties of Ionic Compounds

  • General properties:
    1) Solids with regular shapes due to strong electrostatic forces in the lattice.
    2) High melting points because of strong lattice forces.
    3) Do not conduct electricity in the solid state; conduct when molten or in solution because ions become mobile.
    4) High density due to close packing of ions.
    5) Soluble in water and other polar solvents but insoluble in non-polar organic solvents (e.g., benzene).

  • Solubility in polar solvents is due to attraction between ions and polar molecules.

  • Polar covalent molecules show charge separation; electronegativity is the tendency of an atom to attract bonding electrons toward itself, while electropositivity describes the tendency to push bonding electrons away from itself.

  • Example: Water molecule shows partial charges: ext{δ}^- ext{O} and ext{δ}^+ ext{H} due to higher electronegativity of oxygen relative to hydrogen.

  • Metals are generally malleable and ductile because their bonding is a moving electron cloud; ions can slide past one another without shattering the lattice.

  • Simple molecular structures

    • Consist of discrete molecules held together by relatively weak intermolecular forces (e.g., van der Waals forces) while the atoms within the molecule are held by strong covalent bonds.

    • States: gases, liquids, or solids with low melting points; low density.

    • Examples include: ext{I}2 (iodine), ext{CO}2, ext{NH}3, ext{H}2 ext{O}.

  • Structure of iodine (I2)

    • A strong covalent bond holds iodine atoms in each molecule, while the molecules are held together by weak van der Waals forces.

    • Visual: I–I clusters with weak inter-molecular forces.

Giant Atomic/Molecular Structures

  • These structures consist of atoms or molecules linked by covalent bonds in a three-dimensional network.

  • Diamond

    • Structure: each carbon atom forms covalent bonds with four other carbon atoms in a tetrahedral arrangement; an infinite network, giving a giant 3D covalent structure.

    • Properties:

    • No mobile electrons; cannot conduct electricity.

    • Very hard natural substance; very high melting point.

    • High density (~3.5 g/cm^3).

    • Transparent, sparkling, and shiny.

    • Uses: cutting/drilling (industrial tools), jewelry, and laser applications.

  • Graphite

    • Structure: infinite carbon atoms arranged in layers of hexagonal rings; each carbon is covalently bonded to three others within a layer, forming a giant two-dimensional layer.

    • Layers are held together by weak van der Waals forces, allowing layers to slide over one another (slippery and soft).

    • Within layers, some electrons are mobile, enabling electrical conductivity.

    • Properties:

    • Conducts electricity due to delocalized electrons within layers.

    • High melting point due to strong covalent bonds in layers.

    • Soft and greasy because of layer sliding (slippery).

    • Opaque, dark in color, and shiny.

    • Less resistant to chemical attack than diamond due to open spaces between layers.

    • Density: about 2.3 ext{ g/cm}^3.

    • Uses:

    • Pencil leads (mixed with clay to form leads).

    • Good conductor of electricity; used as electrodes.

    • Lubricants for small bearings (e.g., in dynamos).

    • Brushes for electric motors.

    • Protective coatings on iron (as a protective coating against rust).

Properties of Giant Molecular Structures

  • Typically have high melting and boiling points due to strong bonds.

  • Do not conduct electricity, except graphite which has delocalized electrons.

  • Insoluble in water.

Giant Metallic Structure

  • Structure: giant lattice of positive metal ions surrounded by a sea of delocalized electrons.

  • Properties:

    • High melting and boiling points (except mercury, which is a liquid at room temperature).

    • They conduct electricity in both solid and liquid states due to mobile electrons.

Summary of Key Points and Connections

  • Ionic bonding involves electron transfer and results in an electrostatic lattice that conducts electricity only when ions are mobile (melting/solution).

  • Covalent bonding involves sharing electrons; the number of shared electron pairs determines bond type (single, double, etc.). Lone pairs influence molecular geometry and reactivity.

  • Dative bonding is a covalent bond where both electrons come from the same atom.

  • Metallic bonding features a delocalized electron cloud that enables electrical and thermal conductivity and malleability.

  • Structure follows bonding: ionic bonds lead to giant ionic lattices; covalent bonds give simple molecular or giant covalent networks; metallic bonds give giant metallic structures.

  • Properties broadly align with bonding type: ionic compounds tend to be hard and soluble in polar solvents; simple molecular substances have low mp and poor conductivity; giant covalent networks (diamond) are extremely hard with high mp and non-conductive; graphite conducts electricity due to delocalized electrons; metals conduct due to the electron sea and are malleable.

  • Practical context: understanding bonding explains real-world properties like hardness, conductivity, solubility, and uses (e.g., pencils, lubricants, electronics, cutting tools).

Sample Questions (Study Prompts)

  • Define and differentiate electrovalent (ionic), covalent, coordinate (dative), and metallic bonding. Include formation mechanisms and give named examples.

  • Explain metallic bonding and why metals are good conductors of electricity and heat, and why they can be malleable/ductile.

  • For covalent compounds, draw electronic structures (outer shells only) for select molecules:

    • (i) ext{CCl}4, (ii) ext{PCl}3, (iii) ext{SiH}4, (iv) ext{CHCl}3, (v) ext{PH}3, (vi) ext{CH}2 ext{Cl}_2

  • Using electronic diagrams, show the structure of:

    • (a) calcium atom, (b) chlorine atom, (c) calcium chloride. State differences between electrovalent and covalent compounds.

  • Describe how the structures of the following account for their electrical conduction properties: (a) copper, (b) graphite, (c) diamond.