1. Chemical Bonds Section 1 Introduction to Chemical Bonding

2. Objectives Define chemical bond.
Explain why most atoms form chemical bonds.
Describe ionic and covalent bonding.
Explain why most chemical bonding is neither purely ionic or covalent.
Classify bonding type according to electronegativity differences.

3. Atomic Stability Q. Why do atoms form compounds?
Search for outer shell stability by losing or gaining electrons!
Nobel Gases  8 electrons in os = stable
Ex: Na loses 1 e- to Cl
Chemical Bond  attractive force that holds atoms together in a compound; when atoms gain, lose, or share e-, an attractive force pulls them together to form a compound

5. Oxidation Number  the number that tells you how many electrons an atom has gained, lost or shared to become stable

6. Ion Def. Charged particle that has either gained or loses electrons
RESULT: Now has either more or fewer number of electrons than protons
Ex: sodium fluoride NaF active ingredient in toothpaste  Na loses 1e- to F; Na is +1; F is -1
Ex: potassium iodide KI – ingredient in iodized salt  K loses 1 e -to I; K is +1; I is -1
Cation = positive ion #P > #E
Anion = negative ion #P < #E

7. Ionic Bond
Def. Force of attraction between the opposite charges of the ions in an ionic compound (cations and anions)
Transfer of electrons btwn metals & nonmetals
Ex: magnesium chloride MgCl2  Mg loses 2 electrons to each Cl; Mg is +2; each Cl is -1
RESULT neutral compound = sum of the charges of the ions equals 0
2(Mg) + -1(Cl) + -1(Cl) = 0

8. Covalent Bond
Def. Attraction that forms between atoms when they share electrons
Atoms will be more stable by sharing e- rather than losing or gaining e-

9. Unequal Sharing
Electrons don’t always share equally between atoms in a covalent bond
Strength of attraction of atom to electrons due to:
1. size of atom
2. size of the positive charge in the nucleus (a strong magnet will hold a metal better than a weak magnet)
3. total # of electrons
4. how far are the electrons from the nucleus being shared (a magnet has a stronger pull to a metal when it is next to it rather than a couple inches away)

10. Ex: HCl hydrochloric acid used to clean metal and found in your stomach to digest food
Cl – atoms have a stronger attraction for electrons than H atoms  electrons shared will spend most time near the chlorine atom
RESULT Cl atom has a partial negative charge (Greek delta)  H atom has a partial positive charge
VISUAL: Tug-of-war  stronger team pulls the rope towards them

11. Polar or Nonpolar
Polar molecule  molecule that has a slightly positive end and a slightly negative end although the overall molecule is neutral ex: water
Nonpolar molecule  molecule in which electrons are shared equal in bonds; doesn’t have oppositely charged ends; found in 2 identical atoms or molecules that are symmetric ex: CCl4

14. Ionic or Covalent???
Bonding between atoms of different elements is rarely purely ionic or purely covalent.
Falls somewhere between 2 extremes depending on electronegativity measure of atom’s ability to attract e-

16. Electronegativity Values

17. less than 0.3
greater than 1.7

18. If you still need help understand polarity and electronegativity, watch this video:

19. HOMEWORK
Section Review
pg 177
#1-5

20. Objectives Define molecules and molecular formula.
Explain relationships among potential energy, bond length, and bond energy.
Learn the basic steps used in writing Lewis structures.
Explain how resonance structures are used to represent molecules.

21. Section 2 Covalent Bonding Molecular Compounds
molecule neutral group of atoms that are held together by covalent bonds
Red = O, White = H, Black = C
Molecular compound a chemical compound whose simplest units are molecules
-Many chemical compounds, including most of the chemicals that are in living things and are produced by living things are composed of molecules
-molecule could be single molecule existing on its own or 2 more atoms of the same element or compound

22. chemical formula indicates what elements atoms and numbers of atoms in a chemical compound by using atomic symbols and numerical subscripts
molecular formula shows types and number of atoms combined in a single molecule of a molecular compound

23. Formation of a Covalent Bond
Nature favors chemical bonding most atoms have lower potential energy when bonded to other atoms than as independent atoms.
separated H atoms do not affect each other
PE decreases as atoms are drawn together by attractive forces
PE minimum when attractive forces are balanced by repulsion forces = ideal distance
PE increases when repulsion btwn like charges outweighs attraction between opposite charges

24. Characteristic of the Covalent Bond
Bond length average distance between 2 bonded atoms H-H 75 pm
Form CB= H atoms release energy= amt of energy equals drop in PE
Bond energy energy required to break a chemical bond and form neutral isolated atoms (kJ/mol)

26. Octet Rule
Chemical compounds tend to form so that each atom has an octet of e- in outer energy level by gaining, losing or sharing e-
Draw Fluorine electron configurations:

27. Objectives Learn the basic steps used in writing Lewis structures.
Explain how resonance structures are used to represent molecules.

28. Electron Dot Structure or Lewis Dot Diagram (Gilbert Lewis)
Lewis Dot Diagrams
Electron Dot Structure or Lewis Dot Diagram (Gilbert Lewis)
Def. A notation showing the valence electrons (electrons in outer energy level) surrounding the atomic symbol.

29. Lewis Structures a)H b)P c)Ca d)Ar e)Cl f)Al
1)Write the element symbol.
2)Carbon is in the 4th group, so it has 4 valence electrons.
3)Starting at the right, draw 4 electrons, or dots, counter-clockwise around the element symbol.
On your sheet, try these elements on your own:
a)H b)P c)Ca d)Ar e)Cl f)Al

30. unshared pair- (lone pair) e- not involved in bonding and belong to one atom
Structural formula- shows bonds but not unshared pairs of e- in molecule

31. Single Covalent Bond Made of 2 shared electrons
1 comes from one atom in the bond and 1 comes from the other atom in the bond
Ex: water – O now is stable with 8 e in outer shell and H is stable with 2

32. Multiple Bonds Bonds with multiple pairs of shared electrons
Ex: N2 Nitrogen has 5 e in os and needs 3 to be stable  shares 3 e (triple bond)

34. Lewis Structures for Molecules
Draw the LS for CH3I
1. Write the LS for each atom in the molecule.
2. Determine the total number of valence e- available.
3. Arrange atoms to form skeleton structure of molecule. When present C is central atom. Otherwise, least EN atom is central (Except H).
4. Add unshared pairs of e- to each nonmetal (except H) such that each is surrounded by 8.
5. Count e- to be sure that # of VE=number available. Check to see that atoms have octet.

35. Practice Draw and build the Lewis Structure of NH3
Draw and build the Lewis Structure of H2S
Draw and build the Lewis Structure of SiH4
H-white, N-orange (should only have 3 holes) S-red, Si-black

36. Resonance Structures
def. bonding in molecules or ions that cannot be correctly represented by a single Lewis structure
resonance  aka hybrids constantly alternating from one form to another O3 has a single structure that is avg of 2 structures
use double arrow to indicate resonance

37. Draw the 3 Resonance Structures for SO3

38. Homework
Electron Dot Diagrams and Lewis Structures Worksheet

39. Objectives
Compare and contrast a chemical formula of molecular and ionic compounds
Compare and contrast properties of ionic and molecular compounds
Write Lewis structures for polyatomic ions

40. Section 3 Ionic Bonding and Ionic Compounds
Ionic compound def. composed of + and – ions combined so that # of + and – charges are equal
most exist as crystalline solids= 3D structure of +/- ions attracted to each other
Formula unit def. simplest unit of atoms from which ionic compound can be established ex: NaCl

41. Formation of Ionic Compounds

42. Show formation of … KF, potassium fluoride Na20, sodium oxide

43. Characteristic of Ionic Bonding
*Remember: Nature favors arrangements where potential energy is minimized.
In an ionic crystal, ions minimize PE by combining in orderly arrangement crystal lattice

44. Lattice Energy
def. energy released when one mole of an ionic crystalline compound is formed from gaseous ions negative values mean energy released when crystals are formed

47. Polyatomic Ions def. a charged group of covalently bonded atoms
combine w/ ions of opposite charge to form ionic compounds

48. Metallic Bonding
Occurs bc e- move freely among a metal’s + charged ions e- form a cloud around the metal ion
RESULT: ductility & malleable
Ex: metal hammered into sheets doesn‘t break bc ions are in layers that slide past each other w/out losing their attraction to the
electron cloud
Ex: good conductor of electricity bc outer-level electrons are held weakly
Atoms in metals have a special called metallic bonding

49. Metallic Bonding
Ductility- ability of a substance to be drawn, pulled, or extruded through a small opening to produce a wire
Malleability- ability of a substance to be hammered into thin sheets

50. Chemical Speed Dating Profile
Homework
Chemical Speed Dating Profile
Chapter Review

51. Bell Work Draw the Lewis Dot Diagram for CF4.
Q. What do you think are the limitations of Lewis Structures???
A. They do NOT provide any useful information about the 3D structural arrangement of the molecule.

52. Ch 6 Sec 5 Molecular Geometry
Objectives
Explain VSEPR Theory
Predict shapes of molecules or polyatomic ions using VSEPR theory.
The properties of molecules depend not only on the bonding of atoms but also on the molecular geometry- the 3-D arrangement of a molecule’s atoms in space.
Geometry (from the Ancient Greek geo- "earth", -metron "measurement") is a branch of mathematics concerned with questions of shape, size, relative position of figures, and the properties of space.

53. valence shell electron-pair repulsion
VSEPR Theory
valence shell electron-pair repulsion
states that the repulsion between the sets of valence-level electrons surrounding an atom causes the e- to orient as far apart as possible

54. Let’s look at how electrons in a molecule repel each other and result in different molecular shapes…
BeF2, beryllium fluoride
Draw Lewis Structure:
According to VSEPR theory, the shared pairs will be as far away from each other as possible F atoms are 180° apart= molecule is linear

55. Multiple Bonds?
In VSEPR theory, double and triple bonds are treated the same way as single bonds.

56. Atoms Bonded to Central Atom Lone Pairs of electrons
Molecular Shape
Atoms Bonded to Central Atom
Lone Pairs of electrons
Example
Lewis Structure
Bond Angle(s)
Hybridization 2 1 3 4 5 6

57. Atoms Bonded to Central Atom Lone Pairs of electrons
Molecular Shape
Atoms Bonded to Central Atom
Lone Pairs of electrons
Example
Lewis Structure
Bond Angle(s)
Hybridization
Linear  2
BeF2
180°
Bent  1
SO2 
120° 
H2O
104.5°
Trigonal-Planar  3
BCl3 
Trigonal-Pyramidal
NI3 
109.5° (107)
Tetrahedral  4
SiBr4 
109.5° 
Trigonal-Bipyramidal  5
PCl5 
 90°,120,180
Octahedral 6
SF6 
 90°,
180
Phet interactive stimulation to recognize that molecule geometry is due to repulsions between electron groups.

58. Build Models of Molecules
PROCEDURE 1.Construct Lewis Structure 2. Use materials to indicate the molecular shape of the molecule. 3. Use the protractor to measure the bond angles.

59. CH4
8 Total VE
2. BeH2
4 Total VE
Tetrahedral
Linear

60. Trigonal Planar
Octahedral
3. BF3
24 Total VE
4. SF6
48 Total VE

61. Trigonal-bipyramidal
Bent
5. PCl5
40 Total VE
6. ONF
18 Total VE

62. 7. NH3
8 Total VE
8. H2O
Trigonal-pyramidal
Bent

63. Kahoot Review Game
Homework: Further practice- VSEPR Structures Worksheet *Challenge Structures posted on website

64. Hybridization
VSEPR does not reveal the relationship between a molecule’s geometry and the orbitals of its bonding electrons
Hybridization: the mixing of 2 or more atomic orbitals of similar energy on the same atom to produce a new hybrid atomic orbital of equal energies

65. Methane, CH4 Carbon: 4 VE 2 in 2s orbital, 2 in 2p orbital
When combining with 4 H to make CH4, carbon’s 2s and 2p orbitals hybridize to form 4 new identical sp3 orbitals
*The 3 indicates that 3 p orbitals are in hybridization.
*The sp3 orbital has more energy than the 2s orbital but less than 2p.

66. Hydrogen Bonding

67. London Dispersion Forces