1. Sections 6.1 – 6.4

2. Chemical Bond = A link between atoms Why does it occur? The nucleus of one atom is attracted to the electrons of another.

3. IONIC BOND Ion = Atom which has gained or lost electron(s) Metal = -LEFT side of Periodic Table Weak nucleus / Low Electronegativity -LOSERS of electrons Become positively (+) charged ions (Cations)

4. Nonmetal = -RIGHT side of Periodic Table Strong nucleus / High Electronegativity -GRABBERS of electrons Become negatively (-) charged ions (Anions)

5. Atoms gain or lose valence electrons to become a NOBLE GAS CONFIGURATION Right/Left Side? Metal/ Nonmetal? Lose/Gain?Noble Gas it resembles? Mg Li Cl O

6. Ionic bond = A chemical bond between a cation (+) and an anion (-). Caused by a TRANSFER of electron(s). Usually a metal + a nonmetal

7. COVALENT BOND = A bond caused by a SHARING of electrons Usually a nonmetal + a nonmetal Nonpolar Covalent = Equal sharing of the electrons. Atoms are close in strength Polar Covalent = Unequal sharing of the electrons. One atom is a little bit stronger than the other

8. How do you tell which type of bond it is? -By ELECTRONEGATIVITY A chart of electronegativity will be provided to you. -The greater the difference in electronegativity – the more ionic the bond. -Electrons spend more time closer to the element with higher electronegativity.

9. If the ABSOLUTE VALUE of the electronegativity difference is: GREATER THAN 1.7 = IONIC Bond LESS THAN 0.3 = NONPOLAR COVALENT Bond 0.3 – 1.7 = POLAR COVALENT Bond Examples:

10. METALLIC BOND Usually metals only -The metal gives up valence electrons. -Electrons are free to move about. Atom Electron Sea

11. Covalent Bond = A sharing of electrons Molecule = A group of atoms held by covalent bonds (ex – water) Diatomic Molecule = Molecule with only 2 atoms (7 naturally occurring ones) Molecular Compound = Compound made of molecules Molecular Formula = The type and number of atoms in a molecule (ex – H 2 O)

12. Sharing electrons in a covalent bond makes the atoms more stable and decreases the energy of the atoms. Energy is released when a bond is FORMED. Overlapping of Orbitals – Example H 2 : HH + H2H2

13. Atoms in a compound obtain the electron configuration of a NOBLE GAS to gain stability Exceptions to the Octet Rule:  Hydrogen  Incomplete octets  Expanded octets

14.  Surrounded by only 2 electrons ◦ Can only form single bonds!!

15.  Too few valence electrons  Rather rare: ◦ Boron, aluminum, beryllium  Examples:

16.  Some central atoms form 5 or 6 bonds  Elements with atomic #s greater than 10  Examples:

17. -A picture of the covalent bonds in a molecule -Basic Rules: ◦ Determine a central atom  If C is present – always central  H is never central  Group 17 is never central ◦ Arrange to form skeleton (like a plus sign) ◦ Place dots around each element and connect dots Examples:

18.  More Examples:

19. Single Bond = 1 pair of electrons (2 e - s total) shared between two atoms (longest length; lowest bond energy) Double Bond = 2 pairs of electrons (4 total e - s) shared between two atoms Triple Bond = 3 pairs of electrons (6 total e - s) shared between two atoms (shortest length; highest bond energy)

20.  More Examples:

21.  Same as non-charged molecules, except: ◦ For positive ions – subtract electrons to total # of valence electrons ◦ For negative ions – add electrons to total # of valence electrons  Examples:

22. Ionic Bond = Bond formed by the attraction of a cation to an anion Crystal Lattice = 3-Dimensional network of ions Formula Unit = Simplest ratio of ions NaCl

23. Dot structures for Ionic Compounds: -Want to reach noble gas configuration “Math equation” -Draw an ARROW to show the transfer of e- -Draw as many of each ion as needed Examples:

24. MolecularIonic Bond TypeCovalentIonic StructureIndividual MoleculesCrystal Lattice Strength of BondStrongVERY strong Mp/bpLowHigh DrawingLewis Structures“Math Eqn” Other---Conducts electricity when in water

25.  Do electronegativity difference first!!  Examples:

26. Metals have LOW electronegativity – Will LOSE electrons The steps: -Donate valence electrons to electron sea -Electrons free to move about -All electrons in sea are shared by all atoms

27. Properties of Metals: 1. Good conductors of heat – e- sea shakes 2. Good conductors of electricity – e- in sea can move 3. Malleable – atoms can be pushed closer 4. Ductile – atoms can be pushed closer 5. Luster – light bounces off e- sea

28. Section 6.5

29. VSEPR = Valence Shell Electron Pair Repulsion Theory Valence electrons move as far away from each other as possible 1. Draw Lewis Structure 2. Look at Central Atom 3. Count electron areas (shared areas & lone pairs) 4. Use chart info

30. NameShapeShared AreasLone PairsBond Angles Linear20180° Bent21120° Trigonal Planar 30120° Tetrahedral40109.5° Trigonal pyramidal 31107° Bent22104.5° Trigonal bipyramidal 5090°, 120° Octahedral6090°

31. Examples:

34.  Additional Handout!! Not in note packet!!  Partial Charges ◦ In a polar bond ONLY!! ◦ A tug of war occurs – one atom is “stronger” than the other!  δ + and δ - ◦ Greek letter delta ◦ Compare EN:  Higher EN value of the two = δ -  Lower EN value = δ +

35.  Examples: ◦ H 2 O ◦ CH 4 ◦ NH 3 ◦ CO 2 ◦ And the hardest – CH 3 Cl

36. Mixing a set of atomic orbitals ◦ forms a new set of atomic orbitals with the same total electron capacity ◦ properties and energies intermediate between those of the original unhybridized orbitals. Three types: sp (triple bonds), sp 2 (double bonds), sp 3 (single bonds) Carbon: C BECOMES C 1s 2 2s 2 2p 2 four sp3 hybrid

37. Dipole 1. Dipole = Molecule with overall charge NonPolar With Polar Sites (NPWPS) 2. NonPolar With Polar Sites (NPWPS) = Molecules with area of charge which cancel out Nonpolar 3. Nonpolar = Molecule with no areas of charge

38. How do you tell the difference? -Ask yourself these questions… Is the molecule polar or nonpolar (Difference in EN) Polar (∆EN = 0.3-1.7) Can charge be sliced? YES = DipoleNO = NPWPS Nonpolar (∆EN < 0.3) One straight line so all positive charge is separated from all negative charge; through BONDS only!!

39. AKA – EXTERNAL BONDS BETWEEN The attraction BETWEEN Molecules Types of External Bonds: 1. Dipole-Dipole Interactions -Occur due to attraction between partial charges -Occur between: Two dipoles (strongest) Dipole to NPWPS Two NPWPS (weakest) Hydrogen Bond = Special case of a dipole-dipole external bond that involves a hydrogen atom 

40. 2. London Force -Occurs between nonpolar molecules -Very weak connection The Steps: A.Electrons in one molecule shift instantaneously to one side B.Instantaneous charge results C.Electrons in another molecule are repelled D.Very weak attraction results

41. 1. State of Matter s > l > g **This means that solids have strongest external bonds; gases have weakest bonds 2. Evaporation (Volatility – tendency of a substance to vaporize) slow > fast **Those compounds that evaporate very slow have stronger bonds than those that evaporate quickly 3. Thickness (Viscosity – a measurement of resistance to flow) thick > thin **Substances that are “thicker” have stronger bonds

42. 4. Wetness (Adhesion – the force of attraction) To feel wet the substance must bond to your skin (to the Na+Cl- **If you feel wetness, the substance is bonding to your skin 5. Dissolving LIKE DISSOLVES LIKE **Polar dissolves in polar; nonpolar dissolves in nonpolar