1. Ch. 6 Chemical Bonding
6-1 Introduction to chemical bonding

2. Chemical bond-a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together

3. Types of Bonds
Ionic-results from the electrical attraction between large numbers of cations and anions
Covalent-results from sharing of electron pairs between two atoms

4. Rarely are bonds ever purely ionic or purely covalent but most lie somewhere between

5. fig. 6-3
1. nonpolar covalent bond-the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge
polar- uneven distribution of charge
2. polar-covalent bond-unequal attraction for the shared electrons

6. Fig. 6-2
1. nonpolar-covalent: 0-5% ionic character; electronegativity difference of 0-0.3
2. polar-covalent: 5-50% ionic character; electronegativity difference of
3. ionic: % ionic character; electronegativity difference of >1.7

7. 6-2
6-2 Covalent Bonding and Molecular Compounds

8. Molecule-neutral group of atoms that are held together by covalent bonds
Molecular compound-chemical compound whose simplest units are molecules
Chemical formula-indicates the relative numbers of atoms of each kind in a chemical compound by using symbols and subscripts

9. Molecular formula-the types and numbers of atoms combined in a single molecule of a molecular compound
Diatomic molecule-a molecule containing only two atoms

10. Forming of a bond
The electrons-protons of two atoms are attracted to each other while the electrons-electrons are repelled by each other
The bond forms at a distance at which these two forces become equal, potential energy is at a minimum and a stable molecule forms

11. Covalent Bonds
Bond length-the distance between two bonded atoms at their minimum potential energy, that is, the average distance between two bonded atoms
Bond energy-the energy required to break a chemical bond and form neutral isolated atoms

12. 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
Exceptions: H, B (6), Be(4), P, S, Xe

13. Electron-Dot Notation
An electron-configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element’s symbol
Pg. 170

14. Lewis Structures
Formulas in which atomic symbols represent nuclei and inner-shell electrons, dot-pairs or dashes between two atomic symbols represent electron pairs in covalent bonds, and dots adjacent to only one atomic symbol represent unshared electrons

15. Lone pair (unshared pair)-pair of electrons that belongs exclusively to one atom and not involved in bonding
Structural formula-indicates the kind, number, arrangement, and bonds but not the unshared pairs of the atoms in a molecule ex: F-F

16. Multiple Bonds
single bond-covalent bond produced by the sharing of one pair of electrons between two atoms
Double bond-covalent bond by sharing two pair of electrons
Triple bond-bond by sharing of three pair of electrons

17. Resonance structures
Structure that cannot be correctly represented by a single Lewis structure

18. Ionic Bonding and Compounds
Ionic compound-composed of positive and negative ions that are combined so that the charges are equal
Ions minimize their potential energy (form bonds) by combining in an orderly arrangement known as a crystal lattice
Pg. 177

19. Ionic Bonds/Compounds
Form between a metal (+) and nonmetal (-)
+ and – ions attracted to each other
Hard, brittle
High melting pt/boiling pt
Form crystal lattices
Do not vaporize at room temp
Good electrical conductors

20. Covalent Bonds/Compounds
Share electrons
Between 2 nonmetals
Form molecules
Lower melting pts
Vaporize at room temp

21. Polyatomic Ions Pg. 180 Group of covalently bonded atoms with a charge
NO3-
NH4+

22. 6-4 Metallic Bonding
Metals outer orbital overlap each other so electrons are free to roam from atom to atom
The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons.

23. Metallic Properties
Malleable-ability to be hammered or beaten into thin sheets
Ductile-ability of a substance to be drawn into a wire
Bond strength-the amount of heat required to vaporize the metal is a measure of the bond strength

24. Metallic Bonds Metals bond with other metals
Attraction b/t metals and a surrounding sea of electrons
High electrical and thermal conductivity
Malleable, ductile, and high luster (shiny)
Ability to absorb wide range of light frequencies

25. 6-5 Molecular Geometry
Properties of molecules depend on the bonding and also the geometry or shape.
The polarity depends on the geometry and determines if the molecule is polar

26. VSEPR Theory
States that repulsion of valence electrons surrounding an atom causes them to be oriented as far apart as possible
1. Linear-2 atoms set equal distance apart (180˚) OR 3 atoms because the electrons repel each other
Ex: AB2

27. 2. Trigonal Planar- AB3 difference of 120º
3. Tetrahedral- AB4 difference of 109.5º
VSEPR with unshared electrons
4. Trigonal Pyramidal- AB3E (107º)-lone electrons will also repel each other but shape only involves the atoms involved in the bonding

28. 5. Bent or Angular (105º) AB2E2 or AB2E
6. Trigonal bipyramidal- AB5 ex: PCl5
7. Octahedral- AB6 ex: SF6
Pg. 186 Table 6-5

29. Predict the geometrical shape using the VSEPR theory and by drawing the Lewis Structure for the following:
A. HI E. SO2
B. CBr4 F. Cl4
C. AlBr3 G. BCl3
D. CH2Cl2

30. Using a protractor and pg
Using a protractor and pg. 186, construct the geometrical shapes using gumdrops and toothpicks

31. Intermolecular Forces
Forces of attraction b/t molecules
Higher the boiling pt, the stronger the force
These are generally weaker than bonds that join atoms in molecules
Strongest intermolecular forces are those between polar molecules

32. 1. Dipole-dipole=equal but opposite attractions
dipoles that are additive make it more polar ex: NH3
Dipoles in a molecule that cancel each other make it nonpolar ex: CCl4

33. 2. Induced dipole=temporary dipole develops
Weaker than dipole-dipole force
3. Hydrogen bond=strong type of dipole-dipole between H-F, H-O, H-N
Makes them highly polar
Ex: water H2O

34. 4. London dispersion forces- weak, induced instantaneous dipole created by constant motion of electrons
Only among noble gases and nonpolar molecules due to their low boiling pts.
Their strength increases with the number of electrons in the molecules or with increased mass