1. Ionic Bonding

2. CONCEPTS
Students know atoms combine to form molecules by sharing electrons to form covalent or metallic bonds or by exchanging electrons to form ionic bonds.
Students know salt crystals, such as NaCl, are repeating patterns of positive and negative ions held together by electrostatic attraction.

3. Bonds Forces that hold groups of atoms together and make them function
as a unit.
Ionic bonds – transfer of electrons
Covalent bonds – sharing of electrons

4. The Octet Rule – Ionic Compounds
Ionic compounds form so that each atom, by gaining or losing electrons, has an octet of electrons in its highest occupied energy level.
Metals lose electrons to form positively-charged cations
Nonmetals gains electrons to form negatively-charged anions

5. Ionic Bonding: The Formation of Sodium Chloride
Sodium has 1 valence electron
Chlorine has 7 valence electrons
An electron transferred gives
each an octet
Na: 1s22s22p63s1
Cl: 1s22s22p63s23p5

6. Ionic Bonding: The Formation of Sodium Chloride
This transfer forms ions, each with an octet:
Na+ 1s22s22p6
Cl- 1s22s22p63s23p6

7. Ionic Bonding: The Formation of Sodium Chloride
The resulting ions come together due to electrostatic attraction
(opposites attract):
Na+
Cl-
The net charge on the compound must equal zero

8. Examples of Ionic compounds
Mg2+Cl-2
Magnesium chloride: Magnesium loses two electrons and each chlorine gains one electron
Na+2O2-
Sodium oxide: Each sodium loses one electron and the oxygen gains two electrons
Al3+2S2-3
Aluminum sulfide: Each aluminum loses two electrons (six total) and each sulfur gains two electrons (six total)

9. Metal
Monatomic Cations
Ion name
Lithium
Li+
Sodium
Na+
Potassium K+
Magnesium
Mg2+
Calcium
Ca2+
Barium
Ba2+
Aluminum
Al3+

10. Nonmetal
Monatomic Anions
Ion Name
Fluorine F-
Fluoride
Chlorine
Cl-
Chloride
Bromine
Br-
Bromide
Iodine I-
Iodide
Oxygen
O2-
Oxide
Sulfur
S2-
Sulfide
Nitrogen
N3-
Nitride
Phosphorus
P3-
Phosphide

11. Sodium Chloride Crystal Lattice
Ionic compounds form solid crystals at ordinary temperatures.
Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions.
All salts are ionic compounds and form crystals.

12. Properties of Ionic Compounds
Structure:
Crystalline solids
Melting point:
Generally high
Boiling Point:
Electrical Conductivity:
Excellent conductors, molten and aqueous
Solubility in water:
Generally soluble

13. Metallic Bonding
Strong forces of attraction are responsible for the high melting point of most metals.

14. CONCEPT
Students know atoms combine to form molecules by sharing electrons to form covalent or metallic bonds or by exchanging electrons to form ionic bonds.

15. Metallic Bonding
The chemical bonding that results from the attraction between metal cations and the surrounding sea of electrons
Vacant p and d orbitals in metal's outer energy levels overlap, and allow outer electrons to move freely throughout the metal
Valence electrons do not belong to any one atom

16. Packing in Metals
Model: Packing uniform, hard spheres to best use available space. This is called closest packing. Each atom has 12 nearest neighbors.

17. Metal Alloys
Substitutional Alloy: some metal atoms replaced by others of similar size.

18. Metal Alloys
Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms.

19. Properties of Metals
Metals are good conductors of heat and electricity
Metals are malleable
Metals are ductile
Metals have high tensile strength
Metals have luster

20. Covalent Bonding
Bonding models for methane, CH4. Models are NOT reality. Each has its own strengths and limitations.

21. CONCEPTS
Students know atoms combine to form molecules by sharing electrons to form covalent or metallic bonds or by exchanging electrons to form ionic bonds.
Students know chemical bonds between atoms in molecules such as H2, CH4, NH3, H2CCH2, N2, Cl2, and many large biological molecules are covalent.
Students know how to draw Lewis dot structures.

22. The Octet Rule and Covalent Compounds
Covalent compounds tend to form so that each atom, by sharing electrons, has an octet of electrons in its highest occupied energy level.
Covalent compounds involve atoms of nonmetals only.
The term “molecule” is used exclusively for covalent bonding

23. The Octet Rule: The Diatomic Fluorine Molecule1s 2s 2p
Each has seven valence electrons F 1s 2s 2p F F

24. The Octet Rule: The Diatomic Oxygen Molecule1s 2s 2p
Each has six valence electrons O 1s 2s 2p O O

25. The Octet Rule: The Diatomic Nitrogen Molecule1s 2s 2p
Each has five valence electrons N 1s 2s 2p N N

26. Lewis Structures
Lewis structures show how valence electrons are arranged among atoms in a molecule.
Lewis structures Reflect the central idea that stability of a compound relates to noble gas electron configuration.
Shared electrons pairs are covalent bonds and can be represented by two dots (:) or by a single line ( - )

27. The HONC Rule Hydrogen (and Halogens) form one covalent bond
Oxygen (and sulfur) form two covalent bonds
One double bond, or two single bonds
Nitrogen (and phosphorus) form three covalent bonds
One triple bond, or three single bonds, or one double bond and a single bond
Carbon (and silicon) form four covalent bonds.
Two double bonds, or four single bonds, or a triple and a single, or a double and two singles

28. Completing a Lewis Structure -CH3Cl
Make carbon the central atom (it wants the most bonds, 4)
Add up available valence electrons:
C = 4, H = (3)(1), Cl = 7 Total = 14
Join peripheral atoms
to the central atom with electron pairs. H .. .. .. .. .. H C Cl .. ..
Complete octets on
atoms other than hydrogen with remaining electrons H

29. Bond Length and Bond Energy
Length (pm)
Energy (kJ/mol)
C - C
154
346
C=C
134
612
CC
120
835
C - N
147
305
C=N
132
615
CN
116
887
C - O
143
358
C=O
799
CO
113
1072
N - N
145
180
N=N
125
418
NN
110
942

30. VSEPR and Molecular Geometry
Hemoglobin

31. VSEPR Model (Valence Shell Electron Pair Repulsion)
The structure around a given atom is determined principally by minimizing electron pair repulsions.

32. Predicting a VSEPR Structure
Draw Lewis structure.
Put pairs as far apart as possible.
Determine positions of atoms from the way electron pairs are shared
Determine the name of molecular structure from positions of the atoms.

33. Steric Number 1 1 1 atom bonded to another atom
Steric No.
Basic Geometry
0 lone pair
1 lone pair
2 lone pairs
3 lone pairs
4 lone pairs 1
linear
Steric Number – Number of groups (single bonds,
double bonds, triple bond or unshared electron
pairs attached to the central atom.

34. Steric Number 2
2 atoms, or lone electron pairs, or a combination of the two, bonded to a central atom.
Steric No.
Basic Geometry
0 lone pair
1 lone pair
2 lone pairs
3 lone pairs 2
linear

35. Steric Number 3
3 atoms, or lone electron pairs, or a combination of the two, bonded to a central atom.
Steric No.
Basic Geometry
0 lone pair
1 lone pair
2 lone pairs
3 lone pairs 3
trigonal planar
bent / angular
linear

36. Steric Number 4
4 atoms, or lone electron pairs, or a combination of the two, bonded to a central atom.
Steric No.
Basic Geometry
0 lone pair
1 lone pair
2 lone pairs
3 lone pairs 4
tetrahedral
trigonal pyramid
bent / angular
linear

37. Steric Number 5
5 atoms, or lone electron pairs, or a combination of the two, bonded to a central atom.
Steric No.
Basic Geometry
0 lone pair
1 lone pair
2 lone pairs
3 lone pairs 5
trigonal bipyramid
sawhorse / seesaw
t-shape
linear

38. Steric Number 6
6 atoms, or lone electron pairs, or a combination of the two, bonded to a central atom.
Steric No.
Basic Geometry
0 lone pair
1 lone pair
2 lone pairs
3 lone pairs 6
Octahedral
square pyramid
square planar

39. pentagonal bipyramidal
Steric Number 7
7 atoms, or lone electron pairs, or a combination of the two, bonded to a central atom.
Steric No.
Basic Geometry
0 lone pair
1 lone pair
2 lone pairs
3 lone pairs 7
pentagonal bipyramidal
pentagonal pyramidal