1. How do you think H and O in water are bonded together?
Draw a picture to help your explanation.

2. 6.1 Introduction to Chemical Bonding
Ch. 6 Bonding
6.1 Introduction to Chemical Bonding

3. Chemical Bonds atoms rarely exist alone
when atoms are bonded together, they have less potential energy and are more stable
What is potential energy?
chemical bond – mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together

4. Ionic Bonds
results from electrical attraction between large numbers of cations and anions
atoms donate or accept electrons from each other

5. Covalent Bonds
results from sharing of electron pairs between two atoms
the electrons shared belong to both atoms

6. Ionic vs. Covalent

7. Ionic vs. Covalent
bonding usually does not fall in one category or the other, but somewhere in between
type of bond depends on the elements differences in electronegativities
0.3

8. Practice Determine whether each of the following bonds will be:
Ionic or covalent

9. Patterns What kind of patterns do you see? metals + nonmetals = ionic
nonmetals + nonmetals = covalent

10. Ch. 6 Bonding
6.2 Covalent Bonding

11. What is the difference between covalent and ionic bond?
How do you determine which a compound contains?

12. Molecular Compounds
molecule: neutral group of atoms held together by covalent bonds
molecular compound: compound whose simplest unit is a molecule

13. Formulas
chemical formula: tells the number of each type of atom in a compound
molecular formula: tells the number of each type of atom in a molecular compound
ex. H2O, Cl2, C6H12O2

14. Molecular Compounds
diatomic molecule: a molecules containing only 2 atoms
usually refers to 2 of the same atoms
ex: O2, Br2, F2, etc.
7+1 rule

15. Formation of Covalent Bond

16. Formation of Covalent Bond
two nuclei and two electron clouds repel each other creating an increase in PE
approaching nuclei and electron clouds are attracted to each other to create a decrease in PE

17. Formation of Covalent Bond
a distance between the nuclei is reached in which
repulsion and attraction forces are equal
potential energy is at the lowest point possible
at the bottom of the curve on PE graph

18. Covalent Bonds Bond Length
distance between two bonded atoms at their lowest PE
average distance since there are some vibrations
measured in pm (1012 pm = 1 m)
stronger the bond, shorter the bond

20. Covalent Bonds Bond Energy
energy is released when atoms become because they have lower PE
the same amount of energy must be used to break the bond and form neutral isolated atoms
stronger bond, higher bond energy
average since varies a small amount based on atoms in entire molecule
in kJ/mol

22. Which elements naturally exist as diatomic molecules?
Remember, the rule
How many valence electrons do each of the halogens have?

23. Octet Rule
representative elements can “fill” their outer energy level by sharing electrons in covalent bonds
Octet Rule- a compound tends to form so that each atom has an octet (8) of electrons in its highest energy level by gaining, losing or sharing electrons
Duet Rule- applies to H and He

24. Octet Rule Less than 8: More than 8: Boron: 6 in outer energy level
anything in 3rd period or heavier
because may use the empty d orbital
ex: S, P, I

25. Electron Dot Diagrams a way to show electron configuration
identifies the number and pairing of valence electrons to show how bonding will occur
write the noble gas notation
identify the number of valence
identify how many are paired and how many are alone
do not go by Figure 6-10

26. N Example Nitrogen Sulfur 1s2 2s2 2p3 5 valence 2 are paired
3 are alone
Sulfur
1s2 2s2 2p6 3s2 3p4
6 valence
4 paired (2 pairs)
2 are alone N

27. Lewis Structures like dot diagrams but for entire molecules
atomic symbols represent nucleus and core electrons and dots or dashes represent valence electrons
unshared electrons: (lone pairs) pair of electrons not involved in bonding written around only one symbol
bonding electrons: written in between 2 atoms as a dash

28. Types of Bonds single- sharing of one pair of electrons
weakest, longest
double- sharing of 2 pairs of electrons
stronger and shorter
triple- sharing of 3 pairs of electrons
strongest and shortest
multiple bonds include double and triple bonds

29. Drawing Lewis Structures
find the number of valence electrons in each atom and add them up
draw the atoms next to each other in the way they will bond
add one bonding pair between each connected atoms
add the rest of the electrons until all have 8 (consider exceptions to octet rule)

30. H H C Cl CH3Cl Example 1 methyl chloride C: 4 x 1 = 4 H: 1 x 3 = 3
Cl: 7 x 1 = 7
total = 14 electrons
carbon is central H
H C Cl
duet
octet
duet
octet
duet H
H C Cl

31. H N H H NH3 Example 2 ammonia N: 5 x 1 = 5 H: 1 x 3 = 3 total = 8
N is central
Example 2
H N H H

32. Example 3 N2
nitrogen gas
N: 5 x 2 = 10
10 electrons
N N
N N

33. H C H O Example 4 CH2O formaldehyde C: 4 x 1 = 4 H: 1 x 2 = 2
O: 1 x 6 = 6
total = 12
C is central
H C H O

34. O O O O O O Example 5 O3 ozone O: 6 x 3 = 18 two completely equal
arrangements
the real structure
is an average of these two
where each bond is sharing 3 electrons instead of 4 or 2
O O O
O O O

35. O O O O O O Resonance Structures
resonance – bonding between atoms that cannot be represented in on Lewis structure
show all possible structures with double-ended arrow in between to show that electrons are delocalized
O O O
O O O

36. NO31-
N: 5 x 1 = 5
O: 6 x 3 = 18
total = = 24
Example 6

37. Covalent Network Bonding
a different type of covalent bonding
not specific molecules
lots of nonmetal atoms covalently bonded together in a network in all directions
example:
diamond
silicon dioxide
graphite

38. Ch. 6 Bonding
6.3 Ionic Bonding

39. Ionic Compounds ionic bonds do NOT form molecules
chemical formulas for ionic compounds represent the simplest ratio of ion types
made of anions and cations

40. Ionic Compounds
combined so that amount of positive and negative charge is equal
usually crystalline solid
formula of ionic compound depends of the charges of the ions combined

41. Formation attractive forces: repulsive forces: oppositely charged ions
nuclei and electron clouds of adjacent ions
repulsive forces:
like-charged ions
electrons of adjacent ions

42. Formation
distance between the ions creates a balance between those forces
ions minimize their PE by combining in an orderly arrangement called a crystal lattice

43. Formation specific lattice pattern created depends on: charges of ions
size of ions
Calcium Bromide:
each Ca2+ is surrounded by 8 F-
each F- is surrounded by 4 Ca2+
Sodium Chloride
each Na+ is surrounded by 6 Cl-
each Cl- is surrounded by 6 Na+

44. Ionic vs. Molecular ionic bonds and molecular bonds are both strong
ionic bonds connect all ions together
molecules are more easily pulled apart because intermolecular forces are weak

45. Ionic vs. Molecular Molecular Compounds:
low melting and boiling points
many are gases at room temperature
Because the intermolecular forces of the molecules are weak so they are easily separated

46. Ionic vs. Molecular Ionic Compounds: higher melting and boiling points
all are solid at room temperature
hard: Because of the strong forces, it is difficult for one layer of ions to move past another
brittle: if one layer is moved, the layers come apart completely

47. Ionic vs. Molecular Ionic Compounds: good conductors in liquid state
Because ions are free to move and carry charge
poor conductor in solid state
Because ions are fixed in place

48. Polyatomic Ions
charged group of covalently bonded atoms
Example: CN-

49. NH4+ : ammonium ion

50. O O S O O O H SO42- : sulfate ion OH- : hydroxide ion 5 x 6 = 30
total = = 32
OH- : hydroxide ion
= 8 total O S O O O H

51. Draw the Lewis Structures for
NO21- and PO43-

52. Ch. 6 Bonding
Metallic Bonding

53. Bonding of Metals
the highest energy level for most metal atoms does not contain many electrons
usually have empty p and d block
these vacant overlapping orbitals allow outer electrons to roam freely around the entire metal
the electrons are delocalized – are not with one specific atom

54. Bonding of Metals these roaming electrons form a sea of electrons
around the metal atoms
metal atoms are packed in a crystal lattice
metallic bonding – bonding that results from the attraction between metals atoms and sea of electrons

55. Properties of Metals conductivity luster (shininess)
from the freedom of electrons to move around the atoms
luster (shininess)
contain many orbitals with only small differences in energy
many amounts of energy can be absorbed and emitted

56. Properties of Metals malleability and ductility
bonding is the same in every direction
one layer of atoms can slide past another without friction

57. Bond Strength
depends on the nuclear charge (Z) or the number of protons
depends on the number of electrons in the “sea”
heat of sublimination – amount of heat required to turn solid, bonded metal atoms into gaseous individual atoms

58. Metallic vs. Ionic