1. Modern Chemistry Chapter 6- Chemical Bonding
chemical bond- a mutual electrical attraction between the nuclei and the valence electrons of different atoms that binds the atoms together

2. ionic bond- a chemical bond that results from the electrical attraction between anions and cations

3. Na + Cl  NaCl
sodium loses its one valence electron to form the cation Na+
This allows the atom to become like neon with eight electrons in it outer energy level.
chlorine gains that electron to form the anion Cl-
This allows the atom to become like argon with eight electrons in its outer energy level.
Na+ is attracted to Cl- (opposites attract) and an ionic bond forms

4. covalent bond- a chemical bond that results from the sharing of electron pairs between two atoms
nonpolar covalent bond- a covalent bond in which the bonding electrons are shared equally by the atoms forming the bond
-because their electronegativities are essentially equal
polar covalent bond- a covalent bond in which the bonded atoms have an unequal attraction for the shared electrons
-because one atom has a greater electronegativity than the other

5. Nonpolar covalent bond
H · ·F
Fluorine’s greater electronegativity causes the shared electrons to move closer to it and creates areas of slight positive and negative charge forming a polar covalent bond
Nonpolar covalent bond
H· ·H
Equal electronegativities of the hydrogen atoms cause the pair of electrons to be shared equally and a nonpolar covalent bond to form.

6. Determining Bond Type
Find the absolute difference in the electronegativities of the bonding atoms.
The greater the difference, the greater the
% ionic character which makes it more like an ionic bond.
IF the absolute difference is
< 0.3 the bond is nonpolar covalent
> 0.3 and < 1.7 the bond is polar covalent
> 1.7 the bond is ionic
Do section review problems #1-4 on page 177.

7. Section Review page 177
1- Electron pairs are shared in covalent bonds and electrons are transferred between atoms in ionic bonds. 2-The difference in the electronegativities of bonding atoms determines the bond type. For problem #3 use the electronegativity chart on page a- Li = 1.0 F = = 3.0  ionic bond 3b- Cu = 1.9 S = =0.6  polar covalent 3c- I = 2.5 Br = =0.3  polar covalent 4- c (0.3) < b (0.6) < a (3.0)

8. Covalent Bonding & Molecular Compounds
molecule- a neutral group of atoms that are held together by covalent bonds
molecular compound- a chemical compound whose simplest units are molecules
chemical formula- indicates the relative numbers of atoms of each kind in a chemical compound by using element symbols and numerical subscripts
molecular formula- shows the types and numbers of atoms combined in a single molecule of a molecular compound
diatomic molecule- a molecule containing only two atoms

9. Formation of a Covalent Bond **-see figure 5 on page 179-**
As two atoms come near one another, the nuclei of each atom are attracted to the electrons of the other atom.
This causes a decrease in the potential energy of the atoms.
As a bond between the atoms forms, the potential energy of the system reaches its lowest point. At this point, the two atoms “share” at least one pair of electrons which are then able to move freely between the nuclei of the two atoms in overlapping orbitals.
If the atoms get closer to one another, repulsion between the nuclei increases and the potential energy increases.

10. Bond Length & Bond Energy
bond length- the average distance between two bonded atoms
bond energy- the energy required to break a chemical bond and form neutral, isolated atoms
see figure 7 on page 181
see table 1 on page 182
What is the general relationship between bond length and bond energy (strength) as seen in table 1?

11. The Octet Rule
octet rule- chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its outermost energy level
EXCEPTIONS:
hydrogen atoms are complete with two electrons (H2)
boron atoms are complete with 6 electrons (BF3)
some elements show expanded valence involving “d” orbitals (PF5 & SF6)

12. Electron-Dot Notation
electron-dot notation- is 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 symbol
: F : .

13. Lewis Structures :F-F: ¨ ¨
Lewis structures- are formulas in which atomic symbols represent nuclei and inner shell electrons, dot pairs adjacent to a single atom represent unshared electron pairs, and dashes between the atomic symbols represent covalent bonds between two atoms
¨ ¨
:F-F:
¨ ¨

14. Rules for Drawing Lewis Structures
1- Determine the types and numbers of atoms in the
molecule.
2- Arrange the atoms to form a skeleton structure for the
molecule. If carbon is present, it is the central atom.
Otherwise, the least electronegative atom (except for
hydrogen) is central. Add the electron dot structure for
each atom in the molecule.
3- Connect the atoms with lines to represent covalent bonds
between shared electron pairs.
4- Add unshared pairs of electrons so each atom (other than
hydrogen) has eight electrons.

15. structural formula- indicates the kind, number, arrangement, and bonds but not the unshared pairs of electrons of the atoms in a molecule
F-F

16. Covalent Bonds
single bond- is a covalent bond in which one pair of electrons is shared between two atoms
H· ·H  H-H
-Do practice problems #1-4 on page 186

17. Drawing Lewis Structures
page #1
NH3 N (5 electrons) + 3 H (1 electron each)
H H
H N H  H N H

18. Drawing Lewis Structures
Page #2 H2S
H S  H S
H H

19. Page 186 #3 & 4
3- SiH4 H H Si H H 4- PF3 F P F F

20. multiple covalent bonds- are double (two shared pairs of electrons) or triple (three shared pairs of electrons) bonds
:O=O: :N=N:
¨ ¨
-Do practice problems #1-2 on page 188.
-Do Section Review problems #2, 4, & 5 on page 189.

21. Drawing Lewis Structures
page #1 CO2
O C O O C O

22. Drawing Lewis Structures
page #2 HCN
H C N H C N

23. Section review page 189
2- State the octet rule. Atoms will gain, lose, or share electrons so that each atom has an octet of electrons. 4- a) I Br c) H C C Cl b) H d) Cl H C H Cl Si Cl Br Cl e) F O F

24. Section review page 189
5- H N N H H H H N N H

25. Lewis Structure Practice
Draw the Lewis structures of the following molecular compounds. Also use the model kits to build the molecule. The highlighted formulas represent molecules that contain multiple bonds.
NH3 CO2 N2 O2
HBr H2CO3 C2H6 C2H4
C2H2 PF3

26. NH3  H N H H
HBr  Br H
C2H2  H C C H

27. CO2  O C O
H2CO3  H C O O H O
PF3  F P F F

28. N2  N N
C2H6  H H
H C C H
H H
O2  O O
C2H4  H C C H

29. Drawing Lewis Structures- review
1- Determine which element in the formula will be the central atom(s) of the structure. 2- Make a probable skeleton arrangement of the atoms. 3- Put the correct number of dots to equal the valence electrons of each atom. 4- Draw lines between single electrons of adjacent atoms. 5- If there are extra dots around adjacent atoms, draw multiple bond lines. 6- Make sure each atom has 8 electrons either in unshared pairs or shared bonds. Remember, hydrogen has just two electrons.

30. Lewis Structure Quiz
Draw the Lewis structures for the following compounds. 1- H2O 2- PF3 3- SiO2 4- SeBr2 5- CS2

31. Ionic Bonding & Ionic Compounds
ionic compound- is composed of positive and negative ions that are combined so that the number of positive and negative charges are equal
formula unit- is the simplest collection of atoms from which an ionic compound’s formula can be established
eg. NaCl

32. Formation of Ionic Bonds
An atom of an element with low electronegativity approaches another with high electronegativity.
The highly electronegative atom then transfers an electron from the atom with low electronegativity.
This creates an anion and a cation.
The attraction between the ions forms an ionic bond.
Na Cl  Na Cl -

33. Ionic Crystals
Ionic compounds tend to form an orderly arrangement known as a crystal lattice which then forms crystals.
***see figure 14 on page 191***

34. Comparing Ionic & Molecular Compounds
ionic compound molecular compound
high melting point lower melting point
high boiling point lower boiling point
extreme hardness lower hardness (usually)
brittle less brittle

35. Polyatomic Ions
polyatomic ion- a covalently bonded group of atoms with a positive or a negative charge
Review the list of polyatomic ions given to you by the teacher.

36. Metallic Bonding
metallic bond- a chemical bond resulting from the attraction between metal atoms and the surrounding sea of electrons
The ability of the electrons to move freely between the nuclei of the metal atoms accounts for the unique properties of metals.
This accounts for their being good conductors of heat and electricity.

37. Properties of Metals
luster- the ability to absorb light energy and immediately re-emit it at the same or similar frequency which makes them reflective.
malleable- the ability of a substance to be hammered or beaten into thin sheets
ductility- the ability of a substance to be drawn, pulled, or extruded through a small opening to produce a wire

38. Molecular Geometry
VSEPR theory (valence shell electron pair repulsion)- allows us to predict the shape of molecules. It states that repulsion between the sets of valence level electrons surrounding an atom causes these sets to be oriented as far apart as possible

39. VSEPR Theory Draw the Lewis structure of the compound.
When determining the shape of a molecule using VSEPR, use the following steps:
Draw the Lewis structure of the compound.
Find the central atom(s). Use the letter “A” for this atom.
Count the number of atoms bonded to the central atom. Use the letter ”B” and a subscript for the number of atoms bonded to the central atom “A”.
Count the number of unshared electron pairs around the central “A” atom.
Use the letter “E” and a subscript for the number of unshared electron pairs around the central atom
Use this “ABE” designation to find the molecular shape using table 5 on page 200 of the textbook.
Do practice problems #1 & 2 on page 201.

40. Practice problems page 201
1a- F-S-F S = A F = B = 2 E = 2 AB2E2  bent or angular 1b- Cl-P-Cl P = A B = Cl = 3 E = 1 Cl AB3E  trigonal-pyrimidal

41. H H-C-N-H H H This molecule has two central atoms (C & N) so it has two molecular shapes that are combined.

42. Hybridization
hybridization- is the mixing or two or more atomic orbitals of similar energies on the same atom to produce new hybrid atomic orbitals of equal energies
hybrid orbitals- orbitals of equal energy produced by the combination of two or more orbitals on the same atom
Hybridization explains the unique qualities of a carbon atom with its “sp3” orbitals.

43. Intermolecular Forces
intermolecular forces- the forces of attraction between molecules
dipole-dipole forces- forces of attraction between polar molecules
hydrogen bonding- intermolecular force in which a hydrogen atom bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule
London Dispersion forces- intermolecular attraction resulting from the constant motion of electrons and the creation of instantaneous dipoles

44. Chapter 6 Review
Do the following review problems from pages of the textbook.
#6, 15, 19, 21, 34,
43, & 48

45. End of Chapter Practice
#6 H (2.1) I (2.5)  0.4 = polar covalent
S (2.5) O (3.5) 1.0 = polar covalent
K (0.8) Br (2.8) 2.0 = ionic
Si (1.8) Cl (3.0) 1.2 = polar covalent
K (0.8) Cl (3.0)  2.2 = ionic
Se (2.4) S (2.5)  0.1= nonpolar covalent
C (2.5) H (2.1)  0.4 = polar covalent
#15 H = 1 F = 7 Mg = 2 O = 6
Al = 3 N = 5 C = 4

46. End of Chapter Practice
#19 Li Ca Cl
O C P
Al S

47. End of Chapter Practice
#21 F C F H Se H Br
I N I Br Si Br
I Br Cl
H C H H

48. End of Chapter Practice
#34 AB2 = linear
AB3 = trigonal planar
AB4 = tetrahedral
AB5 = trigonal bipyramidal
AB6 = octahedral

49. End of Chapter Practice
#43 AB3E = trigonal pyramidal
AB2E2 = bent or angular
AB2E = bent or angular

50. Honors Chemistry Chapter 6 Test Review
40 multiple choice questions:
valence electrons, chemical bonds (how they occur)
atoms & potential energy leading to stability and bond formation
polar & nonpolar covalent bonds
difference in electronegativity & % ionic character
using electronegativities, determine if a bond is ionic, polar or nonpolar covalent
definition of molecule, molecular formula (&examples), bond length, octet & octet rule
elements meeting octet rule naturally
how to draw a Lewis structure, identifying a Lewis structure, bonding in Lewis structures

51. Honors Chemistry Chapter 6 Test Review
formula of an ionic compound represents…
lattice energy, crystal lattice
compare properties of ionic & molecular compounds and the strength of their bonds
electrons & charge of polyatomic ions
metallic bonds & their electrons
properties of metals & the cause
properties of ionic crystals
VSEPR definition & use
intermolecular forces (dipole-dipole, hydrogen bonding, London dispersion), their relative strength and properties

52. Chemistry Chapter 6 Test Review
25 multiple choice questions
definitions and functions of valence electrons & chemical bonds
bonding & potential energy
polar & nonpolar covalent bonds & their characteristics
use difference in electronegativity to determine bond type
define molecule, molecular formula, & octet
which elements satisfy the octet rule by themselvs
which elements form multiple covalent bonds
what is necessary to draw a Lewis structure & recognize a correct Lewis structure
properties of ionic vs. covalent compounds
excess (or deficit) electrons in polyatomic ions
valence electrons in metallic bonds
properties of metals and why they occur
definition of VSEPR
intermolecular forces and why they occur, especially dipole-dipole forces
polar molecules

53. VSEPR Lab
H2O H O AB2E2 H
CO2 O C O AB2
H AB4
CH3NH2 H C N H
H H AB3E

54. VSEPR Lab
H2CO H C O AB3 H
CH4 H
H C H AB4

55. VSEPR Lab
C2H6 H H AB4
H C C H
H H AB4
C2H2 H C C H AB2
AB2

56. VSEPR Lab
HCOOH H C O H AB3
AB2E2 O