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Chapter 6

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  • 1. Chapter 6 Chemical Bonding
  • 2.
    • Introduction to chemical bonding
  • 3.
    • What is a chemical bond ???
    • A mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together
  • 4.
    • Why do atoms bond?
    • They are working to achieve more stable arrangements where the bonded atoms will have lower potential energy than they do when existing as individual atoms.
  • 5.
    • Types of Chemical Bonding:
    • 1. Ionic – an electrical attraction that forms between cations (+) and anions (-)
    • 2. Covalent – are formed when electrons are shared between atoms
    • 3. Metallic – formed by many atoms sharing many electrons
  • 6.
    • However….
      • Bonds are never purely covalent or purely ionic.
      • The degree of ionic-ness or covalent-ness depends on property of electronegativity.
  • 7.
    • Recall what electronegativity is:
    • The degree of attraction that an atom has to electrons that are within a bonded compound.
    • (see page 161)
  • 8.
    • To determine the degree of ionic-ness or covalent-ness you must take each of the electronegativities for the elements in the compound and subtract them.
  • 9.
    • If difference is 0-0.3 = nonpolar covalent
    • If difference is 0.3 – 1.7 = polar covalent
    • 1.7 and above = Ionic
  • 10. Ionic/Covalent Character Due to Electronegativity Differences Ionic Polar-Covalent Nonpolar-Covalent 100% 50% 5% 0% 3.3 1.7 0.3 0
  • 11.
    • Sulfur + Hydrogen
    • Sulfur + Cesium
    • Sulfur + Chlorine
    2.5 - 2.1 = 0.4 Polar Covalent 2.5 - 0.7 = 1.8 Ionic 2.5 – 3.0 = 0.5 Polar Covalent The atoms with the larger electronegativity will be the more negative atom!!
  • 12. Practice (p. 177)
    • Use electronegativity differences and Figure 2 to classify bonding between chlorine, Cl, and the following elements: calcium, Ca; oxygen, O; and bromine, Br. Indicate the more negative atom in each pair.
  • 13.
    • In general however…
    • If bonding elements are on opposite sides of the periodic table then they tend to be ionic .
    • If elements are close together, then they tend to be covalent .
  • 14.
    • Covalent Bonding & Molecular Compounds
  • 15.
    • What is a molecule ?
    • A neutral group of atoms that are held together by covalent bonds.
    • May be different atoms such as H 2 O or C 6 H 12 O 6
    • May be the same atoms such as O 2
  • 16.
    • Molecular compounds are made of molecules ….. Not ions!
    • We represent molecular compounds by chemical formulas that show numbers of atoms of each kind of element in the compound. CH 4 - methane
  • 17.
    • Diatomic molecules are those elements that exist in pairs of like atoms that are bonded together.
    • There are 7 diatomic molecules:
    • H 2 N 2 O 2 F 2 Cl 2 I 2 Br 2
  • 18.
    • Octet Rule – Atoms will either gain , lose , or share electrons so that their outer energy levels will contain eight electrons (H is an exception since it can only have 2 in the outer level).
    • These electrons that are being gained, lost, or shared are represented by using the electron dot diagrams .
  • 19. Examples of electron dot notations
    • 1 valence electron
    • 3 valence electrons
    • 5 valence electrons
    • 7 valance electrons
    X X X X
  • 20. Practice (p. 184)
    • Write the electron-dot notation for hydrogen.
    • Write the electron-dot notation for nitrogen.
  • 21. Assignment:
    • Page 177 #3, 4, & 5
    • Page 209 #6 and #19
  • 22. Covalent Bonding
    • Shared electron pairs and unshared pairs:
    • Cl:Cl Shared pair
    • Unshared pairs
  • 23.
    • These electron dot representations are called Lewis structures .
    • Dots represent the valence electrons
  • 24.
    • Lewis structures can also be represented using structural formulas .
    • Dashes indicate bonds of shared electrons (unshared e - are not shown
    • Cl - Cl
    • One pair (2 e - ) is shared here.
  • 25.
    • Lewis structure for ammonia (NH 3 )
  • 26. Example: Draw the Lewis structure of iodomethane, CH 3 I.
    • Step 1:
      • Determine the type and number of atoms in the molecule.
    • Step 2:
      • Write the electron – dot notation for each type of atom in the molecule.
    • Step 3:
      • Determine the total number valence electrons available in the atoms to be combined.
  • 27. Example (con’t)
    • Step 4
      • Arrange the atoms to form a skeleton structure for the molecule. If carbon is present, it is the central atom. Otherwise, the least-electronegative atom is central (except for hydrogen, which is NEVER central). Then connect the atoms by electron-pair bonds.
  • 28. Example (con’t)
    • Step 5
      • Add unshared pairs of electrons to each nonmetal atom (except hydrogen) such that each is surrounded by eight electrons.
    • Step 6
      • Count the electrons in the structure to be sure that the number of valence electrons used equals the number available. Be sure the central atom and other atoms besides hydrogen have an octet.
  • 29.
    • Practice:
      • Draw Lewis structure for methane CH 4
      • Ammonia NH 3
      • Hydrogen Sulfide H 2 S
      • Phosphorus trifluoride PF 3
  • 30.
    • Some atoms can form multiple bonds – especially C, O, & N.
    • Double bonds are bonds that share 2 pair of electrons
    • C=C means C::C
    • Triple bonds share 3 pair
    • C ≡C means C:::C
  • 31. Example (p. 188)
    • Complete just as the other example problems; however, there is an added step.
    • Step 7
      • If too many electrons have been used, subtract one or more lone pairs until the total number of valence electron is correct. Then, move one or more lone electrons pairs to existing bonds between non-hydrogen atoms until the outer shells of all atoms are completely filled.
  • 32. Practice
    • Draw the Lewis structure for carbon dioxide, CO 2 .
    • Draw the Lewis structure for hydrogen cyanide, which contains one hydrogen atom, one carbon atom, and one nitrogen atom.
  • 33.
    • Resonance :
    • Some substances cannot be drawn correctly with Lewis structure diagrams
    • Some electrons share time with other atoms – ex. Ozone – O 3
  • 34.
    • Electrons in ozone may be represented as: O = O–O
    • Other times it may be represented as O–O=O
    • Actually these structures are shared – electrons “resonate” (go back & forth) between them
  • 35.
    • Assignment:
    • p. 189 #4 a – e
  • 36.
    • Ionic Bonding and Ionic Compounds
  • 37. Ionic Bonding & Compounds
    • Ionic compounds are formed of positive and negative ions
    • When combined these charges equal zero
    • Ex: Na = 1+
    • Cl = 1-
    0 charge
  • 38.
    • Ionic substances are usually solids
    • Ionic solids are generally crystalline in shape
    • An ionic compound is a 3-D network of + and – ions that are attracted to each other
  • 39.
    • Crystals in ionic compounds exist in orderly arrangements known as a crystal lattice.
  • 40.
    • Ionic substances are not referred to as “molecules”
    • Ionic substances are referred to as “formula units”
    • A formula unit is the simplest ratio of the ions that are bonded together.
  • 41.
    • The ratio of ions depends on the charges.
    • What would result when F - combines with Ca 2+ ?
            • CaF 2
  • 42.
    • When ions are written using electron dot structures the dots are written and symbols for their charges.
    • Na .  Na +
    • Cl  -
  • 43. Compared to molecular compounds, ionic compounds:
    • Have very strong attractions
    • Are hard, but brittle
    • Have higher melting points and boiling points
    • When dissolved or in the molten state they will conduct electricity
  • 44. Polyatomic Ions:
    • A group of atoms covalently bonded together but with a charge.
    • Sulfate SO 4 2-
    • Carbonate CO 3 2-
    • Nitrate NO 3 -
    • Ammonium NH 4 +
  • 45.
    • Metallic Bonding
  • 46. Metallic Bonding
    • Metals are excellent electrical conductors in the solid state.
    • This is due to highly mobile valence electrons that travel from atom to atom.
    e -
  • 47.
    • Generally metals have either 1 or 2 s electrons
    • p orbitals are vacant
    • Many are filling in the d level
    • Electrons become delocalized and move between atoms
  • 48.
    • A metallic bond is the mutual sharing of many electrons among many atoms.
    • Electrons travel in what is known as the zone of conduction .
  • 49. Metallic Properties
    • High electrical conductivity
    • High thermal conductivity
    • High luster
    • Malleable (can be hammered or pressed into shape)
    • Ductile (capable of being drawn or extruded through small openings to produce a wire)
  • 50. Metallic Bond Strength
    • Varies with nuclear charge and number of electrons shared.
    • High bond strengths result in high heats of vaporization (when metals are changed into gaseous phase)
  • 51.
    • Molecular Geometry
  • 52. Molecular geometry…
    • A molecule’s properties depend on bonding of atoms, but also the molecular geometry .
  • 53. Molecular geometry…
    • Is the three dimensional arrangement of a molecule’s atoms in space.
  • 54. VSEPR Theory
    • Valence Shell Electron Pair Repulsion
    • Electrons around a nucleus repel each other to be as far away from each other as possible.
  • 55. VSEPR Theory
    • AB 2 forms linear molecule as with beryllium
    • However, water (H 2 O) is bent due to electrons repulsion!
  • 56. VSEPR Theory
    • AB 3 forms trigonal planar molecule-ex. ammonia
    • AB 4 forms tetrahedral molecule ex. methane
    • See pg. 200 for other shapes
  • 57. Intermolecular Forces:
    • What happens to liquid molecules when they are heated?
    • As energy is added particles overcome their attraction to each other.
    • IM Forces are the forces of attraction between molecules – not within the molecule.
    • IM forces vary in strength but are weaker than bonds that join atoms
  • 58. Intermolecular Forces:
    • Strongest IM forces exist in polar molecules.
    • Polar molecules act as tiny “dipoles” (equal & opposite charges separated by short distances)
  • 59. Intermolecular Forces:
    • Dipole – dipole forces attract between molecules such as between two water molecules.
    • Positive H region is attracted to negative O region of a different molecule.
  • 60. Intermolecular Forces:
    • Another IM force is Hydrogen bonding .
    • Is a strong type of dipole-dipole force
    • Explains high boiling points of H-containing substances such as water and ammonia
  • 61. Intermolecular Forces:
    • In hydrogen bonding, a hydrogen atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule.
    • The double helix of DNA is held together by hydrogen bonding.
  • 62. Intermolecular Forces:
    • London Dispersion forces :
    • Are very weak bonds
    • Occur due to the fact that since electrons are in constant motion that briefly there are moments where electrons are unevenly distributed and thus the molecule briefly has a charged area.