1. Chemical Bonding

2. Introduction to Bonding

3. Neutral or “free” atoms are rarely found in nature
Most elements exist as part of a compound
Compounds are chemically bonded atoms

4. Chemical bond:
strong attractive force that exists between atoms or ions in a compound.

5. So, why do atoms form bonds?

6. Atoms form bonds to become more stable.

7. Bonding involves only valence electrons

8. Remember…
The noble gases are particularly stable because their outer shell is full of electrons (usually 8)

9. Octet Rule:
Atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons (noble gas configuration)

10. Types of Chemical Bonds
Ionic
Valence electrons are transferred from one atom to another.
Usually formed between a metal and a non-metal.

11. This creates + ions and - ions which are then electrostatically attracted.

12. 2. Covalent Pairs of electrons are shared between atoms.
Usually formed between two non-metals.

13. Two Types of Covalent Bonds
A. Polar Covalent
Electrons are shared unequally.
Usually between two different nonmetals.

14. B. Non-Polar Covalent
Electrons are shared equally.
Usually between atoms of the same element.

15. Metallic
attraction between a metallic cation and delocalized electrons.
Delocalized electrons – valence electrons not held by any specific atom, but free to move from one atom to another.
“Electron Sea”
Usually formed between metals.

16. Bond Properties formula unit (NaCl) molecule (CO2)
ionic vs covalent
formula unit (NaCl) molecule (CO2)
hard & brittle pliable
soluble in H2O PC:
soluble in H2O NPC:
insoluble in H2O

17. Bond Properties ionic vs. covalent
high melting & low melting & boiling points boiling points
electrolytes non-electrolytes
electron Lewis structures transfer diagrams

18. Formula unit
ionic compounds are not found as single molecules.
the simplest ratio of cations to anions
Example: ZnCl2 represents the simplest combination of zinc and chlorine: one Zn 2+ ion and two Cl- ions

19. WATER Contains polar covalent bonds. Called the “universal solvent”
“Like dissolves like”.

20. Properties of Metals Luster High density Good heat conductor
Good electrical conductor
High melting/boiling points
Malleable & Ductile

22. Calculating Bond
Type

23. Very few bonds are purely
one type.
The degree to which bonds
are ionic or covalent can be
estimated by comparing
their electronegativities.

24. Remember:
Electronegativity: the tendency of an atom to attract electrons to itself when chemically bonded to other atoms.

25. Electronegativity Remember:
Increases as you move from left to right across PT
Decreases as you move from top to bottom on periodic table

26. To Calculate Bond Type:
Look up values on Table of Electronegativities.
Subtract.
Locate difference on Bond Type Chart.

27. Bond Type Chart
% ionic character
100% % % %
ionic polar nonpolar
covalent covalent

28. Examples
Ca and Br

29. O and O

30. H and S

32. IONIC BONDS

33. Ionization: formation of an ion by the loss or gain of one or more valence electrons.

34. Cations: positive ions formed by the loss of one or more valence electrons.
Metals tend to form cations.

35. 2) Anions: negative ions formed by the gain of one or more valence electrons
Non-metals tend to form anions.

36. Remember:
Oxidation Numbers tell you how many valence electrons an atom will lose or gain to become stable.

37. We use Electron Transfer Diagrams to represent ionic bonds

38. To write Electron Transfer Diagrams:
Use equation format:
Left of arrow: show electron dot diagrams
Right of arrow: show ions formed and coefficients to give proper ion ratio

39. Stop and do the Examples

40. COVALENT BONDS

41. Covalent bonds are usually formed between non-metals
Remember:
Covalent bonds are usually formed between non-metals

42. We use Lewis Structures to represent covalent bonds

43. Understanding Lewis Structures
Element symbols represent the nuclei and core electrons.

44. Element symbols represent nuclei and core electrons
Dashes between symbols represent shared pairs of electrons
Dot pairs around the outside of a symbol represent unshared electrons

45. Single bond = 1 dash
Double bond = 2 dashes
Triple bond = 3 dashes

46. Central atom is the least electronegative atom.

47. Multiple bonds are most common with
Carbon
Nitrogen
Oxygen

48. To Write Lewis Structures
Count the total number of valence electrons.

49. Choose central atom
(least electronegative)

50. Connect remaining atoms to central atom with single dash.

51. Add remaining electrons around each atom until you have drawn proper number of electrons.

52. Polyatomic ions need to be enclosed in brackets.
The charge should be written in the upper right outside the bracket.

53. Remember Hydrogen can only have two electrons in its valence shell.
Never put extra electrons around Hydrogen!

54. Check for stability.
Each atom should be surrounded by exactly eight electrons.
If all are stable, drawing is complete.
If not, rearrange unshared electrons creating multiple bonds as needed.

55. If not stable, create multiple bonds.

56. and do examples

57. Resonance Structures

58. Some covalent molecules cannot be represented by a single Lewis structure.
The Lewis structure for ozone, O3, can be drawn two ways:

59. Which one is correct?
Neither structure is correct by itself. You will need to draw both.
Use a double headed arrow between the structures to show that the actual molecule is an average of the two possible states.

60. Resonance structure: any one of two or more possible configurations of the same compound that have identical geometry but different arrangements of electrons.

61. Example
Draw the resonance structures for sulfur trioxide.

62. and do homework