The document provides information about ionic bonding, covalent bonding, and molecular structure. It defines ionic bonds as forming between electropositive and electronegative atoms through the transfer of electrons from the metal to the nonmetal. Covalent bonds are described as the sharing of electrons, which can be polar or nonpolar depending on the difference in electronegativity. Molecular structure and the shapes of molecules are explained using valence shell electron pair repulsion theory and hybridization.

1. I N
N D
O
-
+
B +
-
G
+
-
1

2. 2

3. C N
I I
N D -
O N
+
I +
-
+
B O -
G 3

4. IONIC RADII
Zeff
-
+
4

5. MONOATOMIC CATIONS
LOW IONIZATION ENERGY ELEMENTS
ELECTROPOSTIVE ELEMENTS LOSE SOME OR ALL
OF THEIR VALENCE ELECTRONS
⇒LOSE HIGHEST n QUANTUM NUMBER FIRST
TRANSITION ELEMENTS: HIGHEST ns FIRST; THEN d
Na: [Ne]3s1 Na1+: [Ne]
Ca: [Ar]4s2 Ca2+: [Ar]
Fe2+ : [Ar]3d6
Fe: [Ar]3d64s2
Fe3+ : [Ar]3d5
⇒LOSE HIGHEST SUBLEVEL (l ) FIRST
Sn2+: [Kr]4d105s2
Sn: [Kr]4d 5s 5p
10 2 2
Sn4+: [Kr]4d10
REMINDER: CATION RADIUS < ATOM RADIUS
5

6. MONOATOMIC ANIONS
ELECTRONEGATIVE ELEMENTS GAIN ELECTRONS
NUMBER GAINED IS AMOUNT NEEDED TO FILL
VALENCE ORBITALS ....HIGHEST p SUBLEVEL
O: [He]2s22p4 O2-: [He]2s22p6 O2-: [Ne]
I: [Kr]4d105s25p5 I1- : [Kr]4d105s25p6 I1- : [Xe]
TO HELP REMEMBER MONOATOMIC ANION CHARGE:
GROUP NUMBER - EIGHT …….OR
EIGHTEEN - GROUP NUMBER
REMINDER: ANION RADIUS > ATOM RADIUS 6

7. IONIC COMPOUNDS
ELECTROPOSITIVE ELEMENT TRANSFERS OR LOSES
ELECTRON(S) TO THE ELECTRONEGATIVE ELEMENT
METAL LOSES ELECTRON(S) TO NON-METAL
NaF CaF2 FeF3
7

8. 8

9. NDS
BO
ENT
V AL
CO
9

10. TRANSFER ELECTRONS FROM HIGHER E
IONIC ORBITALS TO LOWER E ORBITALS
BONDS: USUALLY: METAL + NON-METAL
NOT INDIVIDUAL MOLECULES
SPHERICAL, NON-DIRECTIONAL CHARGE
NaCl MgF2 BaO
SHARING OF ELECTRONS
COVALENT USUALLY: BETWEEN NON-METALS
BONDS
DIRECTIONAL BONDS
INDIVIDUAL MOLECULES
CO CO2 C2H5OH
10

11. POLAR COVALENT BONDS
MEASURE OF ATOM’S ABILITY TO
ELECTRONEGATIVITY (χ): ATTRACT BONDING ELECTRONS
NON-POLAR COVALENT BOND:∆χ = 0 H2, Cl2, O2
POLAR COVALENT BOND: ∆χ > 0 HCl, H2O, ICl
BOND DIPOLE
δ+
H . ... δ−
Cl BOND POLARITY INCREASES
AS ∆χ INCREASES
χ = 2.1 = 0.9 χ = 3.0
∆χ
ROUGH RULE: ∆χ > 1.8, BOND IS CLASSIFIED AS IONIC
METAL + NON-METAL = IONIC
NON-METAL + NON-METAL = COVALENT
11

12. IN EACH OF THE FOLLOWING, IDENTIFY THE MORE POLAR
BOND AND INDICATE THE DIRECTION OF THE DIPOLE.
REMEMBER: χ INCREASES UP A GROUP
AND ACROSS THE PERIOD
P-F IS MORE POLAR
P-F OR S-F χS > χP
DIPOLE TOWARDS F
∆χ P-F > ∆χ S-F
N-F OR P-F P-F IS MORE POLAR
χN > χP
DIPOLE TOWARDS F
C-H OR O-H χO > χC O-H MORE POLAR IONIC
DIPOLE TOWARDS O
Al-Cl OR Si-Cl χSi > χAl Al-Cl MORE POLAR
DIPOLE TOWARDS Cl
12

13. PICTORIAL REPRESENTATIONS OF
LEWIS VALENCE SHELL ELECTRONS
SYMBOLS:
SHOWS OUTERMOST ELECTRONS IN
4 “ORBITALS” FOLLOWING HUND’S RULE
C DUET
H He
Li Be B C N O F Ne
Na Mg Al Si P S Cl Ar
OCTET
13

14. F F F F F F F F
NONBONDING PAIRS BONDING PAIR
(LONE PAIRS)
TO DETERMINE LEWIS STRUCTURES, YOU NEED:
ELECTRONS REQUIRED:
ER = 8 x NON-H ATOMS + 2 x H ATOMS
VALENCE ELECTRONS AVAILABLE
VE = Σ(VALENCE ELECTRONS IN ALL ATOMS)
SHARED PAIRS (NUMBER OF BONDS) SP = ½ (ER-VE)
LONE PAIRS = ½ (VE) - SP 14

15. GUIDES FOR DETERMINING LEWIS STRUCTURES OF
SYSTEMS OBEYING THE OCTET (DUET) RULE
⇒ DETERMINE ER, VE, SP, & LP
⇒ DRAW MOLECULE WITH SINGLE BONDS
UNLESS NOTED: FIRST ATOM IS CENTRAL
ADD OTHER BONDS TO SATISFY SP
ADD LONE PAIRS TO SATISFY OCTET(S)
ON MOST ELECTRONEGATIVE ATOMS FIRST
WORTH NOTING:
H ATOMS CAN ONLY HAVE 1 BOND (TERMINAL ATOMS)
HALOGENS: 1 BOND UNLESS CENTRAL ATOM
15

16. DRAW LEWIS STRUCTURES FOR:
ER =
CH4 VE =
SP =
LP =
ER =
H2O
VE =
SP =
LP =
16

17. DRAW LEWIS STRUCTURES FOR:
ER =
F2 VE =
SP =
LP =
ER =
HCl
VE =
SP =
LP =
17

18. DRAW LEWIS STRUCTURES FOR:
ER = F P F
PF3 VE =
SP = F
LP =
ER =
NH3
VE =
SP =
LP =
18

19. WHAT IS THE LEWIS STRUCTURE OF O2?
ER = 8 x # O ATOMS
16 ELECTRONS NEEDED
VE = 2 O ATOMS x 6 e- PER ATOM
O O
12 ELECTRONS AVAILABLE
SP = ½(ER-VE) = ½ (16-12) LONE PAIRS = ½ (VE) - SP
2 SHARED PAIRS 2 BONDS ½ (12) - 2 =4
=
WHAT IS THE LEWIS STRUCTURE OF C2H2?
ER = 2 x 8 e- + 2 x 2 e- 20 e-
VE = 2 C ATOMS x 4 e- + 2 H ATOMS x 1 e- 10e-
SP = 1/2 (20 e- - 10e-)
H C C H
5 SHARED PAIRS
19

20. BOND ORDER: NUMBER OF SHARED ELECTRON PAIRS
C C BO = 1 OR SINGLE BOND
C C BO = 2 OR DOUBLE BOND
C C BO = 3 OR TRIPLE BOND
BOND STRENGTH OF A SPECIFIC BOND INCREASES AND
BOND LENGTH DECREASES AS BOND ORDER INCREASES
STRENGTH C O < C O< C O
358 799 1058 kJ/MOLE
LENGTH C O >C O >C O
1.43 1.23 1.13 Ao
20

21. MOLECULAR
STRUCTURE
VSEPR 21

22. VALENCE SHELL ELECTRON PAIR REPULSION
VS E P R
ELECTRON REGIONS OR GROUPS OF NEGATIVE CHARGE
AROUND AN ATOM REPEL ON ANOTHER
ATTAIN POSITION TO MINIMIZE REPULSION
1 ELECTRON GROUP OR ELECTRON REGION IS A:
LONE PAIR
SINGLE BOND
DOUBLE BOND
TRIPLE BOND
THE MOLECULAR SHAPE IS DETERMINED
USING THESE ELECTRON GROUPS
22

23. What does it mean to hybridize?
• Hybridization is the “chemistry word” for
promoting electrons to an empty orbital
36 26
2 4 5 3
1 7 1 7
58 48
The #2 “s” electron becomes a “p” and #5 “p” becomes a “s”!
23

24. Hybridization Summary Table
Total e- Shared Lone
Geometry Hybridization Angle Sketch
pairs pairs pairs
Linear sp 2 2 0
Trigonal
sp2 3 3 0
Planar
Bent sp2 3 2 1
Linear sp2 3 1 2
Tetrahedral sp3 4 4 0
Trigonal sp3
4 3 1
Pyramidal
sp3
Bent Angular 4 2 2
sp3
Linear 4 1 3

25. Hybridization Summary Table
Total e- Shared Lone
Geometry Hybridization Angles Sketch
pairs pairs pairs
Trigonal
sp3d 5 5 0
bipyramidal
See-Saw sp3d 5 4 1
T-shaped sp3d 5 3 2
Linear sp3d 5 2 3
Linear sp3d 5 1 4

26. Hybridization Summary Table
Total e- Shared Lone
Geometry Hybridization Angles Sketch
pairs pairs pairs
sp3d2
ocatahedral 6 6 0
Square sp3d2
6 5 1
pyramidal
Square sp3d2
6 4 2
planar
sp3d2
T-shaped 6 3 3
sp3d2
Linear 6 2 4
sp3d2
Linear 6 1 5

27. THE “MOLECULAR SHAPE” ANALYSIS INCLUDES:
DETERMINE NUMBER OF ELECTRON GROUPS
DETERMINE APPLICABLE SHAPE
NAME THE SHAPE
OBEY OCTET RULE: 2, 3, OR 4 ELECTRON GROUPS
o
180
LINEAR o
120 109
o
TRIGONAL
PLANER TETRAHEDRAL
27

28. WHAT IS THE SHAPE OF THE FOLLOWING:
F P
P F TRIGONAL PYRAMIDAL
F F ~109 o BOND ANGLE
PF3
F
F
H
TETRAHEDRAL
CH4 H C
C H 109 o BOND ANGLE
H
H
HH
O BENT
O
O O OO ~120 o BOND ANGLE
O3
28

29. MORE THAN 4 ELECTRON REGIONS?
5 = TRIGONAL BIPYRAMIDAL
6 = OCTAHEDRAL
29

30. LARGE MOLECULES? DIFFERENT REGIONS!
LINEAR
TRIGONAL
PLANAR
TETRAHEDRAL
CAEFFINE
30

31. BOND ORDER (AGAIN)!
F F
INTERNUCLEAR AXIS σ BOND
ELECTRON DENSITYNODAL PLANE
Π BOND
EVERY BOND CONTAINS 1 σ BOND
MULTIPLE BONDS CONTAIN 1 σ BOND + Π BONDS
SINGLE BOND = σ BOND
F F
DOUBLE BOND = 1 σ BOND + 1 Π BOND
O O
TRIPLE BOND = 1 σ BOND + 2 Π BONDS
H C C H 31

32. HYBRIDIZATION RULES
ONLY FORM IN MOLECULES; NOT ATOMS
ONLY MIX NON-DEGENERATE ORBITALS MOLECULES
COVALENTLY BONDED POLYATOMIC
REQUIRES ENERGY INPUT OR pp
sp, spd BUT NOT ss
H C O
sp2
sp3 1 σ C-C
H C O H
1 σ C-H
O
4 σ BONDS
C-
1 σ C-O
3σ
H O C H OTHER p
O ORBITAL IS
H C O C IN Π BOND
O H
H
32

33. VALENCE BOND THEORY LOCALIZES BONDING ELECTRONS
DIFFICULT TO EXPLAIN RESONANCE
MOLECULAR ORBITAL THEORY DELOCALIZES Π ELECTRONS
C C
FREEDOM TO MOVE
FREEDOM TO SPREAD
33

34. 34

35. L I
A L
E T -
N
+
M +
-
I
C+
N D -
B O 35

36. Metallic Bonding
• With atoms of the same metallic element
• Delocalized electron clouds caused by
metallic atoms being so physically close to
each other
• Known as a “sea of electrons”
• Reason why metals are such good
conductors
36