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Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
Chemical Bonding
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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.

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