2. CHEMICAL BONDS

3. Chemical Bond  Mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together.  Ionic Bond Bonding that results from the electrical attraction between large numbers of cations and anions  A large difference in electronegativity between two atoms in a bond will result in ionic bonding

4.  Covalent (Molecular) Bond Sharing electrons pairs between two atoms A small difference in electronegativity between two atoms in a bond will result in covalent bonding.  Electronegativity A measure of the ability of an atom in a chemical compound to attract electrons.

5. Valence Electrons  Electrons in the highest occupied energy level of an elements’ atom  Determines the chemical properties of an element  Examples:

6. The Octet Rule  Atoms tend to gain, lose, or share enough electrons to become surrounded by eight valence electrons  Non-metallic elements gain electrons or share them  Metallic elements lose electrons

7. Electron Dot Notation  An electron configuration notation in which only the valence electrons of an atom of a particular element are shown  Indicated by dots placed around the element’s symbol ○ Example:

8. Learning Check A. X would be the electron dot formula for 1) Na 2) K3) Al B. X would be the electron dot formula 1) B 2) N3) P

9.  Dot notation can be used to represent molecules Example: (H : H) represents a shared electron pair  An unshared pair or lone pair is a pair of electrons that is not involved in bonding and belongs exclusively to one atom Example: Lone Pair

10. Lewis Structures  Electron distribution is depicted with Lewis electron dot structures  Electrons are distributed as shared or Bond Pairs or unshared or Lone Pairs

11. Steps: Write electron dot notation for each atom in molecule. Determine total number of valence electrons. Arrange atoms to form skeleton structure for molecule. If Carbon is present it will go in the center. Otherwise least electronegative atoms will be in the center. Hydrogen never is in the center

12. Steps Continued Connect atoms by electron-pair bonds Add unshared pairs of electrons so each nonmetal is surrounded by 8 electrons Count the electrons to see it matches the number of valence electrons If too many electrons create double or triple bonds.

13. Lewis Structure Bond Formula To determine how many bonds exist in a molecule, use the following formula: N-A = # of Bonds 2 Electrons are shared and represented by a dash lone electrons are represented by dots.

14. N-A = # of Bonds 2  Where:  N = # of needed electrons,  ( 8 for all elements but H, which is 2.)  A = # of available electrons (the number of valence electrons.)

15. The least electronegative element is the central atom.

16.  Example:  Write the Lewis structure of NH 3  The total number of valence electrons is:  The number of electrons needed is: A = 8 1 x 8 = 8 3 x 2 = 6 N = 14

17. The Skeleton structure is: N – A 2 14 – 8 2 = 3 Bonds

18.  Connect the atoms with electron pairs. Remember 8 electrons are needed to obey the octet rule  Finish the structure by using the remaining electrons as lone pairs  Check that the final Lewis structure has the correct number of valence electrons (8) except Hydrogen (2)

19. Write the Lewis structure of H 2 CO 2 (H)2 x 1= 2 1 (C)1 x 4= 4 1 (O)1 x 6= 6 A = 12 Needed (N) Available (A) 2 (H)2 x 2 1 (C)1 x 8 1 (O)1 x 8 = 4 = 8 N = 20

20. Write the Lewis structure of H 2 CO cont. H | C= O | H N – A 2 20 – 12 2 = 4 Bonds

21. Cl This is the chlorine molecule, Cl 2 Single Covalent Bond Covalent bond produced by the sharing of one pair of electrons between two atoms or

22.  Multiple Covalent Bonds  Double Bond Covalent bond produced by the sharing of two pairs of electrons between two atoms ○ Shown by either two side-by-side pairs of dots or by two parallel dashes

23. O O = For convenience, the double bond can be shown as two dashes. O O

24. Triple Bond Covalent Bond produced by sharing of three pairs of electrons between two atoms Carbon forms a number of compounds containing triple bonds Lewis structures for molecules that contain carbon, oxygen or nitrogen, remember that multiple bonds between pairs are possible

25.  If too many electrons have been used, subtract one or more pairs until the total number of valence electrons is correct. Then move one or more lone electron pairs to existing bonds between non-hydrogen atoms until the outer shells of all atoms are completely filled.

26. Ionic Bonding & Ionic Compounds  Ionic Compound Composed of positive (cations) and negative (anions) ions that are combined so that the numbers of positive and negative charges are equal.

27.  Most ionic compounds exist as crystalline solids. A crystal of any ionic compound is a 3-D network of positive and negative ions mutually attracted to each other. Formation of an ionic bond can be viewed as a transfer of electrons

29. Ionic Bonds: One Big Greedy Thief Dog!

30. Crystal Lattice In an ionic compound, the ions minimize their potential energy by combining in an orderly arrangement. The distance between ions and their arrangement in a crystal represents a balance among all forces.

31. Properties of Ionic Compounds  High melting points  Solids at room temperature  Soluble in polar solvents, insoluble in Nonpolar solvents  Molten compounds & aqueous solutions conduct electricity

32. Metallic Bonding  Chemical bonding is different in metals than it is in ionic, molecular or covalent network compounds

33. Metallic Bonding Chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons An attraction of the free- floating valence electrons for the positively charged metal ions.

34. Metallic Properties Good conductors of electricity and heat Malleability ○ Ability of a substance to be hammered or beaten into thin sheets Ductility ○ Ability of a substance to be drawn, pulled, or extruded through a small opening to produce a wire.

35. The Properties of Molecular (Covalent) Compounds

36. Intermolecular Forces  Reviewing what we know  Low density  Highly compressible  Fill container Solids High density Slightly compressible Rigid (keeps its shape) Gases

37.  Intermolecular forces – occur between molecules  Intramolecular forces – occur inside the molecules

38. Polar Molecule vs. Nonpolar Molecule One very important property of molecule is whether it is polar or nonpolar. If the electrons in a molecule are not evenly distributed, the molecule can have a negative and positive side. A polar molecule will dissolve in another polar substance (such as water). Nonpolar molecule will dissolve in another nonpolar substance (such as carbon tetrachloride).

39. Polar vs. Nonpolar Rules 1.If the central atom has no lone pairs and has all the same types of atoms attached to it, then the molecule is nonpolar. 2.If the central atom has no lone pairs but different atoms attached to it, the molecule is polar. 3.If the central atom has lone pairs, the molecule is polar.

40. Dipole – dipole attraction  A dipole is created by an uneven charge distribution (electronegativity)  Dipole-Dipole is an electrostatic attraction between polar molecules.

41. Hydrogen Bonding  Occurs between H and highly electronegative atom (for example N, O, F)

42. Hydrogen Bonding  Affects physical properties Boiling point

43. London Dispersion Forces  Attraction of instantaneous and induced dipoles; exist between all molecules.  Formation of instantaneous dipoles

44. London Dispersion Forces  Nonpolar molecules

45. London Dispersion Forces  Become stronger as the sizes of atoms or molecules increase

46. Valence Shell Electron Pair Repulsion Theory Planar triangular Octahedral Trigonal bipyramidal

47. VSEPR theory  (Valence Shell Electron Pair Repulsion) states that repulsion between valence electrons causes these sets to be oriented as far apart as possible.

48. Molecular Shape Atoms bonded to central atom (B) Lone pairs (E) Type of Molecule Linear 20AB 2 Bent or angular 21AB 2 E Trigonal- planar 30AB 3

49. Molecular Shape Atoms bonded to central atom Lone pairsType of Molecule Tetrahedral40AB 4 Trigonal- pyramidal 31AB 3 E

50. Molecular Shape Atoms bonded to central atom Lone pairsType of Molecule Bent or angular 22AB 2 E 2 Trigonal- bipyramidal 50AB 5 Octahedral60AB 6