1. Chapter 6: Ionic Bonds and Some Main-Group Chemistry
1/2/2019
Copyright © 2008 Pearson Prentice Hall, Inc.

2. Ionic Bonding Occurs in ionic compound
Results from transferring electron
Created a strong attraction among the closely pack compound Na
Na+ Cl
Cl-

3. Ions and their configurationMg
Mg2+ O
O2-
Atoms
Ions
Fe:
[Ar] 4s2 3d6
- 2 e-
- 3 e-
Fe2+:
Fe3+:
[Ar] 3d6
[Ar] 3d5
Mn:
Mn2+

4. Chapter 6: Ionic Bonds and Some Main-Group Chemistry
1/2/2019
Ionic Radii
Effective nuclear charge.
Copyright © 2008 Pearson Prentice Hall, Inc.

5. Chapter 6: Ionic Bonds and Some Main-Group Chemistry
Ionization Energy
1/2/2019
Ionization Energy (Ei): The amount of energy necessary to remove the highest-energy electron from an isolated neutral atom in the gaseous state.
Students sometimes mistakenly equate ionization energy to the amount of energy required to form an ion.
Increasing Ei
Decreasing Ei
Copyright © 2008 Pearson Prentice Hall, Inc.

6. Chapter 6: Ionic Bonds and Some Main-Group Chemistry
Ionization Energy
1/2/2019
Boron has a lower Ei due to a smaller Zeff
(shielding by the 2s electrons)
Boron has a lower Ei due to a smaller Zeff
(shielding by the 2s electrons)
Copyright © 2008 Pearson Prentice Hall, Inc.

7. Chapter 6: Ionic Bonds and Some Main-Group Chemistry
1/2/2019
Electron Affinity
Electron Affinity (Eea): The energy released when a neutral atom gains an electron to form an anion.
Books use different conventions. The one this book uses assigns a negative sign to electron affinity.
Copyright © 2008 Pearson Prentice Hall, Inc.

8. Ionic Bonds and the Formation of Ionic Solids
Chapter 6: Ionic Bonds and Some Main-Group Chemistry
Ionic Bonds and the Formation of Ionic Solids
1/2/2019
1s2 2s2 2p63s1
1s2 2s2 2p6 3s23p5 Na + Cl
Na1+
Cl1-
1s2 2s2 2p6
1s2 2s2 2p6 3s2 3p6
A transfer of electrons from the metal to the nonmetal.
Copyright © 2008 Pearson Prentice Hall, Inc.

9. Chapter 6: Ionic Bonds and Some Main-Group Chemistry
1/2/2019
The Octet Rule
Octet Rule: Main-group elements tend to undergo reactions that leave them with eight outer-shell electrons.
Copyright © 2008 Pearson Prentice Hall, Inc.

10. Chapter 6: Ionic Bonds and Some Main-Group Chemistry
1/2/2019
The Octet Rule
This is important in Lewis structures and molecular shapes.
Copyright © 2008 Pearson Prentice Hall, Inc.

11. Chapter 7
Covalent Bond and Molecular Structure

12. Types of Chemical Bonds
Bonds: a force that holds groups of two or more atoms together and makes them function as a unit
Required 2 e- to make a bond
Bond energy: amount of energy required to form or to break the bond

13. Chapter 7: Covalent Bonds and Molecular Structure
1/2/2019
The Covalent Bond
Covalent Bond: A bond that results from the sharing of electrons between atoms.
Copyright © 2008 Pearson Prentice Hall, Inc.

14. Polar Covalent Bonds: Electronegativity
Chapter 7: Covalent Bonds and Molecular Structure
1/2/2019
Polar Covalent Bonds: Electronegativity
Depending on the relative electronegativities of the two atoms sharing electrons, there may be partial transfer of electron density from one atom to the other. When the electronegativities are not equal, electrons are not shared equally and partial ionic charges develop.
Atoms with greater electronegativities will attract more of the shared electron density to themselves, causing a “polarity” to the bond
Electronegativity is a measure of the ability of an atom in a molecule to draw bonding electrons density to itself.
Copyright © 2008 Pearson Prentice Hall, Inc.

15. Polar Covalent Bonds: Electronegativity
Chapter 7: Covalent Bonds and Molecular Structure
Polar Covalent Bonds: Electronegativity
1/2/2019
NaCl
Cl2
NaCl
HCl
Copyright © 2008 Pearson Prentice Hall, Inc.

16. Relationship Between Electronegativity and Bond Type
Predicting bond polarity
Atoms with similar electronegativity (Δ EN <0.4) –form nonpolar bond
Atoms whose electronegativity differ by more than two (Δ EN > 2) – form ionic bonds
Atoms whose electronegativity differ by less than two (Δ EN < 2) – form polar covalent bonds

17. Polarity
Polar covalent bonds – the bonding electrons are attracted somewhat more strongly by one atom in a bond
Electrons are not completely transferred
More electronegative atom: δ- . (δ represents the partial negative charge formed)
Less electronegative atom: δ+
Example: For each of the following pairs of bonds, choose the bond that will be more polar
a. H-P, H-C b. N-O, S-O

18. Lewis Structures or Lewis Formula
represents how an atom’s valence electrons are distributed in a molecule
Show the bonding involves (the maximum bonds can be made)
Try to achieve the noble gas configuration
The common pattern

19. Examples Draw Dot Lewis structure for the following atoms: Na Mg C SCo

20. Rules for multiple atoms
Duet Rule: sharing of 2 electrons
E.g H2
H : H
Octet Rule: sharing of 8 electrons
Carbon, oxygen, nitrogen and fluorine always obey this rule in a stable molecule
E.g F2, O2
Bonding pair: two of which are shared with other atoms
Lone pair or nonbonding pair: those that are not used for bonding

21. Electron-Dot Structures
Chapter 7: Covalent Bonds and Molecular Structure
1/2/2019
Electron-Dot Structures
Think of this section as an introduction. It is much easier to write electron-dot structures using the rules listed in the next section.
Copyright © 2008 Pearson Prentice Hall, Inc.

22. Rules for Wring Dot Lewis structure
Step 1: Calculate the total number of valence electrons of all atoms in the molecule
Step 2: Create a skeletal structure using the following rules:
Hydrogen atoms (if present) are always on the “outside” of the structure. They form only one bond
The central atom is usually least electronegative. It is also often unique (i.e,. the only one atom of the element in the molecule). Remember, there might be no “central” atom.
Connect bonded atoms by line (2-electron, covalent bonds
Step 3: Place lone pairs around outer atoms (except hydrogen) so that each atom has an octet
Step 4: Calculate the number of electrons you haven’t used. Subtract the number of electrons used so far, including electrons in lone pair and bonding pairs, from the total in Step 1. Assign any remaining electrons to the central atom as lone pair
Step 5: If the central atom is B (boron) or Be (beryllium), skip this step
If the central atom has an octet after step 4, skip this step
If the central atom has only 6 electrons, move a lone pair from an outer atom to form a double bond between outer atom and the central atom
If the central atom has only 4 electrons, do Step 5a to two different outer atoms (i.e, form two double bonds) or twice to one outer atom (i.e., form one triple bond)

24. Examples Give the Lewis structure for the following HBr CO2 BH3 OF2
NH4+ NO3-,

25. Resonance
when there is more than one Lewis structure for a molecule that differ only in the position of the electrons, they are called resonance structures
the actual molecule is a combination of the resonance forms – a resonance hybrid
it does not resonate between the two forms, though we often draw it that way
look for multiple bonds or lone pairs
Rules for drawing resonance
Resonance structures must have the same connectivity
only electron positions can change
Resonance structures must have the same number of electrons
Second row elements have a maximum of 8 electrons
bonding and nonbonding
third row can have expanded octet
Formal charges must total same
Better structures have fewer formal charges
Better structures have smaller formal charges

26. Complete the three resonance formulas below by using all correct arrows pushing. The draw a resonance hybrid