1. Metals  Lose e-s  positive ion
Octet Rule & Ions
Atoms may or may not be stable due to their nuclear composition
They become more stable by undergoing radioactive decay
Atoms may or may not be stable due to their electronic composition
They can become more stable by undergoing ion formation
Metals  Lose e-s  positive ion
called cations Na+ Mg2+ Fe3+
Ions are electrically charged particles
Nonmetals  Gain e-s  negative ion
called anions Cl - O2- P3-

2. The most stable atoms are the Noble gases
Atoms lose or gain e-s to acquire a noble gas electron configuration (8 valence electrons) is called the octet rule.
Exception [He] 2 valence e-s.

3. Na Na+ + e- Metals lose electrons due to their low ionization energy.
IE = energy required to remove an electron
M + energy → M+ + e-
Na Na+ + e-
[Ne] = Stable State valence e-s

4. Nonmetals gain electrons due to their high electron affinity.
EA = energy change when an electron is added.
X + e− → X− + energy
[Ar] = stable state valence e-s
Cl + e- Cl-

5. Most common ionic charges
Transition metals
Charge is always given by means of:
The name Iron (III) chloride
The formula FeCl3

6. Ionic Compounds

7. Na2S From charges to chemical formulas Mg202 Mg0
When ionic particles react, they do so in a ratio that cancels their charge.
This is what is referred to as “charge balance”.
Na2S
“cross-the-charges” Al Al203
Always keep the simplest ratio Mg
Mg Mg0

8. Practice
1. How many e-s will the following elements need to lose or gain to fulfill the octet rule? N Cl Na Al O Mg
2. What is the ionic charge developed by the following elements?
N Cl Na Al O Mg
3. The symbol for the ion formed by :
N Cl Na Al O Mg
4. Write the formula of the ionic compound formed by:
a. Calcium and chlorine b. Magnesium and nitrogen c. Barium and oxygen

9. Naming & Writing Ionic Formulas
NaCl Sodium chloride
FeCl3 Iron (III) chloride

10. Polyatomic Ions Note: A polyatomic ion taken more than once
Compounds containing polyatomic ions are named as ionic compounds. Na+ SO Na2S04 Sodium sulfate
Al3+ SO Al2 (S04)3 Aluminum sulfate
Figure 05-T07
Title:
Names and Formulas of Some Common Polyatomic Ions
Caption:
Note:
A polyatomic ion
taken more than once
use parenthesis

12. Practice 5. Provide the name for the following ionic compounds:
Mg3N Fe2O3
Cu2S AuCl3
CaSO Cu(NO2)2
6. Provide the formula for the following ionic compounds:
iron(III) chloride chromium(III) oxide
aluminum bicarbonate calcium hydroxide
7. Provide the formula for a compound containing ammonium ion(s) and phosphate ion(s).
8. Draw the electron-dot structure for SO2

13. Between two nonmetal atoms.
Covalent Compounds
Covalent bonds form:
Between two nonmetal atoms.
When atoms share valence electrons to complete octets (duet for hydrogen).
Denoted by a line
Gilbert Newton Lewis
Figure UN
Title:
Formation of a Hydrogen Molecule
Caption:
In the covalent bond in H2, the shared electrons give the noble gas configuration of He to each of the H atoms. Thus the atoms bonded in H2 are more stable than two individual H atoms.
Has a noble gas electron count
Becomes more stable
Has lower energy

14. H H An electron pair being shared is denoted by a line.
In his honor, we refer to electron-dot structures as Lewis structures.
H H
Lewis further proposed that by counting valence electrons, it would be possible to predict how many bonds an atom can form:
H O N C

15. Electron-dot structure = Lewis structure
All valence electrons are shown as dots
or bonding electrons are shown as lines.
F2 Fluorine
In a single bond
One pair of electrons is shared.
Carbon dioxide C : O .. C : O ..
In a double bond,
Two pairs of electrons are shared.
Hydrogen cyanide
In a triple bond.
Three pairs of electrons are shared. : N C H : N C H

16. Steps Guidelines to Writing Lewis Structures NH3 N 5 H  1 x 3 = 3
1. Use the group number to determine the total number of valence e-s
N 5
H  1 x 3 = 3
Total = 8
2. Draw a “skeleton” by connecting the atoms with lines.
First element of the formula is at the center, the others are surrounding. H |
H—N—H
3. Provide octets or duets to surrounding atoms & find the remaining e-s. H |
H—N—H 8 -6 =2
4. Use the remaining electrons to provide an octet to the central atom.
5. If not enough e-s remain, pull surrounding e-s to form double or triple bonds. H |
H—N—H
6. Any e-s remaining after all octets are provided, are given to the central atom.

17. SO3 S 6 O  6 x 3 = 18 Total = 24 O | O—S—O 24 O -24 | =0 O—S—O O ||
1. Use the group number to determine the total number of valence e-s
S 6
O  6 x 3 = 18
Total = 24
2. Draw a “skeleton” by connecting the atoms with lines.
First element of the formula is at the center, the others are surrounding. O |
O—S—O
3. Provide octets or duets to surrounding atoms & find the remaining e-s.
4. Use the remaining electrons to provide an octet to the central atom. O |
O—S—O 24
-24 =0
5. If not enough e-s remain, pull surrounding e-s to form double or triple bonds. O ||
O—S—O
6. Any e-s remaining after all octets are provided, are given to the central atom.

18. Naming & Writing Covalent Formulas
STEP 1 Name the first nonmetal by its element name.
STEP 2 Name the second nonmetal with an --ide ending.
STEP 3 Add prefixes to indicate the number (from subscripts) of atoms of each nonmetal. Mono is always omitted for the first element and usually omitted for the second one.
Exception is CO carbon monoxide
Examples of compounds of N and O:
NO nitrogen oxide nitric oxide nitrogen (I) oxide
NO2 nitrogen dioxide
N2O dinitrogen oxide
N2O3 dinitrogen trioxide
N2O4 dinitrogen tetroxide nitrous oxide nitrogen (II) oxide
N2O5 dinitrogen pentoxide

19. Chapter 6, Figure 6.4, A flowchart illustrates naming for ionic and molecular compounds.

20. Practice
9. Write the electron-dot formula for PCl3, phosphorus trichloride.
10. Write the electron-dot formula for CO2, carbon dioxide.
11. What is the name for P4S3
12. Write the correct formula for each of the following:
A. phosphorus pentachloride B. dinitrogen trioxide
C. sulfur hexafluoride
13. Identify each compound as ionic or covalent and give its correct name.
A. SO B. BaCl2
C. (NH4)3PO D. Cu2CO3
E. N2O4

21. Electronegativity & Bond Polarity
Electrons can be Equally shared
Electrons can be Unequally shared +1
+17
Yes, separation of charge = polar
+1 -1
= 0
-1 +1
= 0
No separation of charge = nonpolar

22. Electronegativity indicates the attraction of an atom for shared electrons.
Linus Pauling
When a separation of charges exists in a bond, it is a polar bond.
Always on the highest EN value

23. Chapter 5, Table 5.14

24. Practice 14. Which element is most electronegative?
a. N b. F c. O d. C
15. Use the electronegativity difference to identify the type of bond between the following:
nonpolar covalent (NP), polar covalent (P), or ionic (I).
A. K-N B. N-O
C. Cl-Cl D. H-Cl
16. Which compound contains a polar covalent bond?
a. HBr b. NaF c. F2 d. H2
17. Which of the following would have the most polar bond?
a. Cl – Cl b. F – Cl c. F – I d. I - Br

25. Shapes & Polarity of Molecules
On paper, molecules appear to be flat, two dimensional objects.
Instead, they have a 3-D structure
The valence-shell electron-pair repulsion theory (VSEPR) predicts the shape based on bonding and lone pairs (“electron groups”) around the central atom.
4 Electron Groups  Four bonding
4 Electron groups  Two Bonding and Two Lone Pairs
VSPER pronounced V (like in Victor) – and (spr)

26. Chapter 5, Table 5.16

27. Polar bonds can produce  a polar molecule  or a nonpolar molecule
Based on EN, a bond can be polar.
Based on shape, a bond can be polar, but a molecule nonpolar

28. For multiple dipoles, it helps to use “center of (+)” and “center of (-)”

29. Example Determine the polarity of the H2O molecule.
Solution: Write the electron-dot structure
Determine the shape & bond polarity.
Determine the molecular polarity. 29 29

30. Practice
18. What is the shape of the SO3 molecule. Provide a drawing of the molecule.
a. Trigonal planar b. Bent c. Tetrahedral d. Trigonal Pyramidal
19. What is the shape of the NH3 molecule. Provide a drawing of the molecule.
a. Linear b. Bent c. Tetrahedral d. Trigonal pyramidal
20. Determine the shape of each of the following molecules and whether they are polar or nonpolar. If polar show the direction of the dipole.
PBr3 B. HBr
C. Br2 D. SiBr4