1. Chapter 15/16
Bonding

2. Types of Chemical Bonds
A. Ionic bond:
Results from the electrostatic attraction between positive and negative ions
B. Covalent bond:
Resulting from the sharing of electrons

3. A+ : B-
1. A + B
A : B
2. Chemical bonds between unlike atoms are never really ionic or covalent, fall somewhere in between

4. 3. How do we determine whether a bond is ionic or covalent. a
3. How do we determine whether a bond is ionic or covalent? a. Compare electronegativities (ability of an atom to attract electrons) b. Electronegativity difference

5. 4.% ionic character 100% 50% 5% 0% ionic Nonpolar covalent
Polar covalent
ionic
Nonpolar covalent

6. 5. Nonpolar covalent: a. share electrons equally b. No charge 6
5. Nonpolar covalent: a. share electrons equally b. No charge 6. Polar covalent: a. share electrons unequally b. Have a partial positive (δ+) and a partial negative charge (δ-)

7. 7. Examples: Using the following chart classify bonds between the following, which bond will be more negative? Element EN value Sulfur 2.5 Hydrogen 2.1 Cesium 0.7 Chlorine 2.5

8. a) H-S b) Ce-S c) S-Cl
Metallic bonding
1. happens with solids and liquids
2. metal atoms give electrons
3. electrons slide through the material (not transferred to a negative)

9.  Covalent bonding and molecular compounds
1. Share electrons
2. Still using valance electrons
3. Share 1 pair of electrons: a single bond

10. 4. Two shared pairs: a double bond
5. Three shared pairs: a triple bond
6. Examples:
a) H H + H = H:H
H-H

11. b. O2
c. N2

12. 7. Octet rule: Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level a. He is happy with 2 electrons b. exceptions to the octet rule

13. 1. Group IIIA metals get 6 electrons when covalently bound
2. Group IIA metals get electrons when covalently bound

14. 3. Group IA metals get 2 electrons when covalently bound
4. Elements in periods 3-7 can have more than 8 electrons in their octet

15. Structures
1. Molecular structure: shows the types and numbers of atoms
PH3
2. Lewis structure:
a. Atomic symbols represent nuclei and inner-shell electrons
b. Dot-pairs or dashes between two atomic symbols

16. c. Will have shared pairs and unshared pairs ex. Cl-Cl 3
c. Will have shared pairs and unshared pairs ex. Cl-Cl 3. Structural formula a. Indicates the kind, number, arrangement, and bonds of the atoms in a molecule

17. b. H-P-H H Examples: a. Draw the Lewis structure of iodemethane (CH3I)
b. Ammonia (NH3)

18.  4 Step Method
1. Add total valence e-
2. Draw skeleton structure
3. Subtract bonded e- from total valence e-
4. Distribute remaining e- (in pairs) to satisfy octet rule

19. Examples:
SO2
NO3-

20. Resonance
1. Refers to bonding in molecules that cannot be correctly represented by a single Lewis structure
2. SO3

21. Ionic bonding and Ionic Compounds
1. Electron-dot structure
a. NaCl
b. CaF2

22. 3. PCl3
*4. AlCl3

23. 2. Lattice: The energy released when one. mole of an ionic crystalline
2. Lattice: The energy released when one mole of an ionic crystalline compound is formed from gaseous ions.
Metallic Bonding
1. metallic bond: resulting from the attraction between positive ions and surrounding mobile electrons

24.  The VSEPR Theory
(Valence-Shell Electron Pair Repulsion Theory)
The pairs of electrons in the valence shell will repel each other until they are as far away from other electron pairs as they can be.

25. Linear NA NA AB HCl H – Cl Linear 2 0 AB2 CO2 O = C = O
Molecular Atoms bonded Lone pairs Type Formula Formula
Shape to central atom of electrons of molecule example Structure
around central
atom
Linear NA NA AB HCl H – Cl
Linear AB CO O = C = O
Bent or AB2E H2O O
H H
Trigonal AB AlCl Cl
Planar Al
Cl Cl
Trigonal AB3E NH N
Pyramidal H H H

26. E: stands for unshared electron pair
Molecular Atoms bonded Lone pairs Type Formula Formula
Shape to central atom of electrons of molecule example Structure
around central
atom
Tetrahedral AB CH H C
H H H
E: stands for unshared electron pair
***Make sure you put dots around the elements where they needed them.

27. 2. How do we use this table?
A. Draw the shape of PCl3.
(Is it eitherAB3 or AB3E)

28. B. CS2 Is it AB2 or AB2E  To move or not to move: 1. Polarity
a. Nonpolar molecules
1. electrons shared equally
2. molecule doesn’t move

29. 1. electrons shared unequally
b. Polar molecules
1. electrons shared unequally
2. molecule would move in the direction of strongest pull

30. 3. dipole: equal but opposite charges separated by a short distance
a. Direction of dipole points to the negative pole
b. ✚
c. Examples:
1. H2O O-H
1.4 polar bond

31. 2. CF4
3. CSO

32. Bond angles
1. Linear
2. a. Bent (2 unshared pairs) o
b. Bent (1unshared pairs) o
3. Triangular planar
4. Tetrahedral
5. Triangular pyramidal
6. Triangular bipyramidal , 1200
7. Octahedral o

33. Hybridization
1. The blending of two or more orbitals to make a new orbital with a new shape
2. Helps explain the bonding and geometry of some Group 15 and Group 16 elements

34. 3. Hybridization rules How to find the (-) areas: a
3. Hybridization rules How to find the (-) areas: a. each bond type counts as 1 area of – charge b. Each unshared pair of electrons count c. NEVER COUNT HYDROGEN

35. d. areas of hybridization (-) charge types 2 sp 3 sp2 4 sp3

36. e. Examples 1. CH4 2. NH3 3. CO2 4. NOF

37.  Bond length
A. The average distance between two bonded atoms, minimum potential energy
1. Longest - single bonds
2. Middle - double bonds
3. Shortest - triple bonds

38. B. Bond energy 1. Energy required to break a chemical bond and form neutral atoms 2. can be related to bond length a. Long bond length needs low energy to break it apart b. Short bond length will need high energy

39. 3. chart bond Bond length Bond energy H-H 74 436 C-H 110 414 N-H 98
389
N-N
140
159

40.  Other properties of polarity
1. Boiling point
a. Temp at which a liquid turns to a gas
b. The greater attraction of charges, the higher the boiling point

41. c. Ionic - highest
NaCl + -
- +
d. Polar covalent - middle
&+ &-
&- &+
e. Nonpolar covalent - lowest

42. 2. Intermolecular forces
a. Dipole-dipole
1. Polar molecules
2. The attraction between the pole of one molecule and the - pole of a different molecule
3. Has a middle boiling point

43. b. London dispersion force
1. Nonpolar
2. The attraction between the temporary pole of one molecule with an adjacent molecule
3. Low boiling point

44. c. hydrogen bonding 1. attraction between a hydrogen atom and an unshared pair of electrons on a strongly electronegative atom (fluorine, oxygen, nitrogen 2. EN difference between hydrogen and fluorine makes them highly polar

45. 3. highest boiling points
Melting points/freezing points
A. ionic bonds
1. Strongest attraction
2. Highest melting/freezing point
B. Polar bonds
1. Somewhat strong bonds
2. Middle values

46. C. Nonpolar bonds
1. Weakest attractions
2. Lowest melting/freezing points
D. Ex. Rank from lowest to highest
1) boiling point
2) freezing /melting points
H2O NaCl F2

47. Quiz Q’s
Draw the correct shape of compound.
Name the shape.
What are the bond angles?
What is the hybridization?
What type of molecule is it (polar or nonpolar)?
BCl3