Chemical Bonding

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  • 1. Chapter 6 CHEMICAL BONDING
  • 2.
    • Mutual electrical attraction between the nuclei and valence electrons of different atoms that finds that atom together
    • Why are most atoms chemically bonded together?
      • Atoms are less stable existing by themselves than when they are combined
      • By bonding with each other, atoms decrease in potential energy, which means they create more stable arrangements of matter
    CHEMICAL BOND
  • 3. CHEMICAL BONDING VIDEO
  • 4.
    • Ionic bonding
      • Results from the electrical attraction between cations and anions
      • Atoms completely give up electrons to other atoms
    • Covalent Boning
      • Results from the sharing of electron pairs between two atoms
      • The shared electrons are “owned” equally by the two bonded atoms
    TYPES OF BONDS
  • 5. COVALENT BONDS VIDEO CLIP
  • 6. IONIC BOND VIDEO CLIP
  • 7.
    • Non-polar
      • The bonding electrons are shared equally by the bonded atoms
      • Results in a balanced distribution of electrical charge
      • Example: hydrogen-hydrogen bond
    • Polar
      • There is an uneven distribution of charge
      • The bonded atoms have an unequal attraction for the shared electrons.
    TYPES OF COVALENT BONDS
  • 8. POLAR VS NON-POLAR COVALENT BONDS
  • 9.
    • Molecule
      • Neutral group of atoms that are held together by covalent bonds
    • Molecular Compound
      • Chemical compound whose simplest units are molecules
    • Chemical Formula
      • Indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts
    • Molecular Formula
      • Shows the types and numbers of atoms combined in a single molecule of a molecular compound
    • Diatomic Molecule
      • Molecule that contains only two atoms
    MOLECULAR COMPOUNDS
  • 10.
    • Bond length
      • The average distance between two bonded atoms
    • Bond Energy
      • The energy required to break a chemical bond and form neutral isolated atoms
      • Measure in kilojoules per mole (kJ/mol)
      • Example:
        • 436 kJ/mol of energy is needed to break hydrogen-hydrogen bonds in one mole of hydrogen molecules
    CHARACTERISTICS OF COVALENT BONDS
  • 11.
    • Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level
    • Exceptions to the rule:
      • Hydrogen
        • Forms bonds in which it is only surrounded by two electrons
      • Boron
        • Forms bonds in which it is surrounded by six electrons
    OCTET RULE
  • 12.
    • Electron configuration notation in which only the valence electrons of an atom of a particular element are shown
    • Valence electrons are indicated by dots placed around the element’s symbol
    ELECTRON DOT NOTATION
  • 13.
    • Formulas in which atomic symbols represent nuclei and inner-shell electrons
    • Dot pairs or dashes between two atomic symbols represent electron pairs in covalent bonds
    • Dots that are adjacent to only one atomic symbol represent unshared electrons
    • Unshared pair (lone pair) – a pair of electrons
    • that is not involved in bonding and that
    • belongs exclusively to one atom.
    LEWIS STRUCTURES
  • 14.
    • Indicates the kind, number, arrangement, and bonds but not the unshared pairs of atoms in a molecule
    STRUCTURAL FORMULAS
  • 15.
    • Single
      • Covalent bond in which one pair of electrons is shared between two atoms
    • Double
      • Covalent bond in which two pairs of electrons are shared between two atoms
    • Triple
      • Covalent bond in which three pairs of electrons are shared between two atoms
    • Multiple bonds
      • Either double or triple bonds
    BONDS
  • 16. Multiple Covalent Bonds
    • Double Bond – a covalent bond produced by the sharing of two pairs of electrons between two atoms.
    Triple Bond – a covalent bond produced by the sharing of three pairs of electrons between two atoms. N N O O N N O O
  • 17. IONIC BONDING AND IONIC COMPOUNDS
    • Ionic Compound – a compound composed of positive and negative ions that are combined so that the number of positive and negative charges are equal.
    • Formula Unit – the simplest collection of atoms from which an ionic compound’s formula can be established.
    • Ex. NaCl
  • 18. FORMATION OF IONIC COMPOUNDS Na + + Cl - NaCl Ca 2+ + F - CaF 2 Ca 2+ + N 3- Ca 3 N 2
  • 19. CHARACTERISTICS OF IONIC BONDING
    • Crystal lattice – an orderly arrangement of ions.
    • Lattice energy – the energy released when one mole of an ionic crystalline compound is formed for gaseous ions.
    • Ionic bonds are stronger than molecular bonds.
  • 20. POLYATOMIC IONS
    • Polyatomic ions – a charged group of covalently bonded atoms.
  • 21. METALLIC BONDING
    • Within a metal, the vacant orbitals in the atoms’ outer energy level overlap, allowing outer electrons of atoms to roam freely throughout the entire metal.
    • Delocalized electrons – electrons that do not belong to one atom, but can freely move about the metal’s network of empty atomic orbitals.
    • Metallic Bonding – the chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons.
  • 22.
    • The shiny appearance of metals are due to the absorption of a high range of light frequencies, resulting in exciting/de-exciting electrons.
    • Malleability – the ability of a substance to be hammered or beaten into thin sheets.
    • Ductility – the ability of a substance to be drawn, pulled, or extruded through small opening to produce a wire.
    • Heat of Vaporization – the amount of heat required to vaporize a metal, which is the measure of the strength of the bonds that hold metal together.
  • 23. MOLECULAR GEOMETRY
    • Molecular Geometry – the three-dimensional arrangement of molecule’s atoms in space.
    • Molecular Polarity – the uneven distribution of molecular charge.
  • 24. VSEPR THEORY
    • Valence Shell Electron Pair Repulsion (VSEPR Theory) – the repulsion between the sets of valence-level electrons surrounding an atom cause these sets to be oriented as far apart as possible.
  • 25.
    • Example: CH 4 (bonding pairs only, no lone pairs)
    Key: Must consider both bonding and lone pairs in minimizing electron repulsion. Lewis Structure VSEPR Structure
  • 26. • Example: NH 3 (both bonding and lone pairs). Lewis Structure VSEPR Structure Molecular Shape
  • 27. VSEPR APPLICATIONS The previous examples illustrate the strategy for applying VSEPR to predict molecular structure: 1. Construct the Lewis Dot Structure 2. Arrange the bonding and lone electron pairs in space such that repulsions are minimized .
  • 28. Case: Linear Structure ( AX 2 ): angle between bonds is 180° Example: BeF 2 180°
  • 29. Case: Trigonal Planar Structure ( AX 3 ): The angle between bonds is 120° Example: BF 3 120°
  • 30. Case: Pyramidal ( AX 3 E ): Bond angles are <120° structure is nonplanar due to repulsion of lone-pair. Example: NH 3 107° VSEPR Structure Molecular shape Lewis
  • 31. Case: Tetrahedral ( AX 4 ): the angle between bonds is ~109.5° Example: CH 4 109.5°
  • 32. Note: for ‘Tetrahedral’, the actual angle may vary slightly from 109.5°, due to size differences between bonding and lone pair electron densities bonding pair: more elongated, less repulsive lone pair: puffier, more repulsive
  • 33. Example of distorted tetrahedron: water ( AX 2 E 2 ): the angle is reduced to 104.5° by repulsion of the lone pairs “ bent” VSEPR structure Molecular shape
  • 34. Case: Trigonal Bipyramidal ( AX 5 ): non-equivalent bond positions: three in-plane (equatorial, 120°), and two at 90° to plane (axial) Example, PCl 5 90° 120°
  • 35. Octahedral ( AX 6 ): all angles are 90° Example SF 6 90° Lewis VSEPR
  • 36. INTERMOLECULAR FORCES
    • Intermolecular Forces – the force of attraction between molecules.
    • Dipole – equal and opposite charges that are separated by a short distance.
    • Dipole-Dipole Forces – forces of attraction between polar molecules.
    H - Cl
  • 37.
    • Hydrogen Bonding – the intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electromagnetic atom in the nearby molecule.
    • London Dispersion Forces – the intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles.
      • Act between all atoms and molecules
      • The only intermolecular forces acting among noble-gas atoms, nonpolar molecules, and slightly polar molecules
    • Only intermolecular among noble gasses and non-polar molecules.