1. Chapter 12
Chemical bonding

2. 12.1 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

3. Ionic Bonding Occurs in ionic compound
Results from transferring electron
Created a strong attraction among the closely pack compound

4. Covalent Bonding Formation of a covalent Bond
Two atoms come close together, and electrostatic interactions begin to develop
Two nuclei repel each other; electrons repel each other
Each nucleus attracts to electrons; electrons attract both nuclei
Attractive forces > repulsive forces; then covalent bond is formed

5. 12.2 Electronegativity
Electronegativity (EN): the ability of an atom in a molecule to attract the shared electron in a bond
Metallic elements – low electronegativities
Halogens and other elements in upper right-hand corner of periodic table – high electronegativity

6. 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+

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

8. Examples
Using the electronegativity values given in Figure 12.3, arrange the following bonds in order of increasing polarity: H-H, O-H, Cl-H and F-H
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
c. O-F, O-I d. N-H, Si-H

9. 12.3 Polarity and Dipole Moment
a vector quantity from the center of the positive charge to the center of negative charge
Represents with an arrow
E.g Draw the dipole moment for HF, H2O, HCl, OF

10. 13.4 Stable Electron Configurations and Charges on Ions
Atoms in stable compounds almost always have a noble gas electron configuration
Predicting Formulas of Ionic Compound
Electrons lost by a metal come from the highest-energy occupied orbital
Electrons gained by a nonmetal go into lowest-energy unoccupied orbital

11. Examples
Predicting formulas of Ionic compound by showing how they loses or gains electrons
Ca and O
Li and S

12. 12.5 Ionic Compound
Lattice energy (U) – the sum of the electrostatic interaction energies between ions in a solid
Refer to the breakup of a crystal into individual ions

13. 12.6 Lewis Structures
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

14. Rules 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

15. Lewis Structures B C N O F use common bonding patterns
C = 4 bonds & 0 lone pairs, N = 3 bonds & 1 lone pair, O= 2 bonds & 2 lone pairs, H and halogen = 1 bond, Be = 2 bonds & 0 lone pairs, B = 3 bonds & 0 lone pairs
often Lewis structures with line bonds have the lone pairs left off
their presence is assumed from common bonding patterns
structures which result in bonding patterns different from common have formal charges B C N O F
Tro, Chemistry: A Molecular Approach

16. Rules 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

17. Rules
Place lone pairs around outer atoms (except hydrogen) so that each atom has an octet
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

18. Rules
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)

19. Examples
Draws a dot Lewis structure for
Na, O, Cl, H2O, HCl, CCl4

20. 12.7 Lewis Structures of Molecules with Multiple Bonds
Recall: Elements typically obey the octet rule; they are surrounded by eight electrons
single bond: involves two atoms sharing one electron
Double bond: involves two atoms sharing two pair of electrons
Use 6N + 2 Rule
N = number of atoms other than Hydrogen

21. Molecular Structure: The VSEPR Model
VSEPR: Valence shell electron pair repulsion
Predicting the molecular structure of molecules formed from nonmetals
Bonding and nonbonding pairs (lone pairs) around given atom are positioned as far as possible

23. Examples Give the Lewis structure for the following CF4, CO, SO42-
NH4+

24. Vector Addition
Tro, Chemistry: A Molecular Approach

25. Tro, Chemistry: A Molecular Approach