1. Additional Chapter 5 information
Write on a piece of paper for notes:
Ion – an atom or group of bonded atoms that has a positive or negative charge.
(not neutral anymore)
Ionization – Any process that results in the formation of an ion
A + energy  A+ + e- (atom loses an e- and becomes more
positive)
A + e-  A- + energy (atom gains an e- and becomes more
negative)

2. Ionic Radii
Cation – A positive Ion. Formed by the loss of one or more electrons. Always leads to a decrease in atomic radius.
KITTY …. KITTY …. KITTY
 Anion – A negative Ion. Formed by adding one or more electrons. Always leads to an increase in atomic radius.

3. Comparing Cations and Anions
Section 3 Electron Configuration and Periodic Properties
Chapter 5
Comparing Cations and Anions
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4. Valence Electrons – electrons available to be lost, gained, or shared in the formation of chemical compounds.
these electrons are in the highest energy level and are often located in incompletely filled main-energy levels.
For Main-Group elements, the valence e- are the electrons in the outermost s and p sublevels
*Table 4 p. 160

5. Section 3 Electron Configuration and Periodic Properties
Chapter 5
Valence Electrons
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Visual Concept

6. Chap 6: Introduction to Chemical Bonding
Learning Targets:
1) Define chemical bond.
2) Explain why most atoms form chemical bonds.
3) Describe ionic and covalent bonding.
4) Explain why most chemical bonding is neither
purely ionic nor purely covalent.
6) Classify bonding type according to electronegativity differences.

7. Introduction to Chemical Bonding
Chemical Bond – is a mutual electrical attraction between the nucleus and valence electrons of different atoms that binds the atoms together.
By themselves, most atoms have a high potential energy.
Nature favors arrangements in which potential energy is minimized.
By bonding with each other, atoms decrease in potential energy, creating more stable arrangements of matter.

8. Chemical Bond Chapter 6 Section 1 Introduction to Chemical Bonding
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Visual Concept

9. Types of Chemical Bonding
Ionic Bonding – chemical bonding that results from the electrical attraction between large numbers of cations and anions. (METALS)
atoms completely give up electrons to other atoms (fig. 1 page 176)
Covalent Bonding – results from the sharing of electron pairs between two atoms (fig. 1 p. 176) (NONMETALS)
Nonpolar-covalent bond – bonding electrons are shared equally by bonded atoms, resulting in a balanced distribution of electrical charge.
(Fig. 3 p. 177)
Polar-covalent bond bonding electrons are NOT
shared equally by bonded atoms resulting in
an unbalanced distribution of electrical charge.
The Polar bear DOES NOT share! (Fig. 3 p. 177)

10. Ionic Bonding Chapter 6 Section 1 Introduction to Chemical Bonding
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Visual Concept

11. Covalent Bonding Chapter 6 Section 1 Introduction to Chemical Bonding
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Visual Concept
Comparing ionic and covalent bonds

12. Comparing Polar and Nonpolar Covalent Bonds
Section 1 Introduction to Chemical Bonding
Chapter 6
Comparing Polar and Nonpolar Covalent Bonds
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Visual Concept

13. - Bonding between atoms is never purely ionic and is rarely purely covalent.
- Electronegativity – measure of an atom’s ability to attract electrons. (chap. 5 p. 161)
- The degree to which bonding between atoms of two elements is ionic or covalent can be estimated by calculating the difference in the elements’ electronegativities.
sample problem A p. 177 (3 slides)

14. - Fig. 2 p. 176 – shows
Electronegativity from 0 – 0.3 = nonpolar-covalent bond (0 – 5% ionic)
Electronegativity from 0.3 – 1.7 = polar –covalent bond (5 – 50% ionic)
Electronegativity from 1.7 – 3.3 = Ionic bonds (50 – 100% ionic)
- Bonding between two atoms of the same element is completely COVALENT (called a diatomic element)
- Carbon is unique because it can form 4 covalent bonds and bond with other carbon atoms to form long chain molecules of different sizes and shapes.

15. Using Electronegativity Difference to Classify Bonding
Section 1 Introduction to Chemical Bonding
Chapter 6
Using Electronegativity Difference to Classify Bonding
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Visual Concept

16. Chemical Bonding, continued
Section 1 Introduction to Chemical Bonding
Chapter 6
Chemical Bonding, continued
Sample Problem A
Use electronegativity values listed in Figure 20 from the previous chapter in your book, on page 161, and Figure 2 in your book, on page 176, to classify bonding between sulfur, S, and the following elements: hydrogen, H; cesium, Cs; and chlorine, Cl. In each pair, which atom will be more negative?

17. Chemical Bonding, continued
Section 1 Introduction to Chemical Bonding
Chapter 6
Chemical Bonding, continued
Sample Problem A Solution
The electronegativity of sulfur is 2.5. The electronegativities of hydrogen, cesium, and chlorine are 2.1, 0.7, and 3.0, respectively. In each pair, the atom with the larger electronegativity will be the more-negative atom.
Bonding between Electroneg. More-neg-
sulfur and difference Bond type ative atom
hydrogen 2.5 – 2.1 = 0.4 polar-covalent sulfur
cesium – 0.7 = 1.8 ionic sulfur
chlorine – 2.5 = 0.5 polar-covalent chlorine

18. Section 2: Covalent Bonding and Molecular Compounds
Learning Targets:
Define molecule and molecular formula.
2) Explain the relationships among potential energy, distance between approaching atoms, bond length, and bond energy.
3) State the octet rule.
4) Construct basic Lewis structures.
5) Explain how to determine Lewis structures for molecules containing single bonds, multiple bonds, or both.
6) Understand that molecules, with covalent bonds are always NONMETALS.

19. Covalent Bonding and Molecular Compounds
Molecule – a neutral group of atoms that are held together by covalent bond nonmetals
could be two or more of the same atom (O2) or two or more different atoms (C6H12O6) (nonmetals)
Diatomic Molecule – molecule containing only two atoms (O2 ; CO)
A molecule is the smallest possible amount of a compound

20. Visual Concepts
Chapter 6
Molecule

21. Comparing Monatomic, Diatomic, and Polyatomic Molecules
Visual Concepts
Chapter 6
Comparing Monatomic, Diatomic, and Polyatomic Molecules

22. 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 atomic symbols and numerical subscripts
Molecular Formula – shows the types and numbers of atoms combined in a single molecule of a molecular compound. (H2O) (nonmetals)

23. Chemical Formula Chapter 6
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Chemical Formula
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Visual Concept

24. Formation of a Covalent Bond
Fig. 5 p. 179
A) Separated Hydrogen atoms each have high Potential Energy (PE)
B) PE decreases as atoms are drawn together by attractive forces.(Protons are attracted to electrons)
There is also Repulsive forces (P vs P and e- vs e-)
C) PE is at a minimum when
the attractive forces are
balanced with the
repulsive forces.
D) PE increases when repulsion between like charges is greater than the attraction between opposite charges.

26. Formation of a Covalent Bond
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Formation of a Covalent Bond

27. Formation of a Covalent Bond
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Formation of a Covalent Bond

28. Characteristics of the Covalent Bond
Bond Length – the distance between 2 bonded atoms at their minimum PE
Bond Energy – the energy required to break a chemical bond and form neutral isolated atoms. Measured in Kilojoules per mole.
Bond Length is inversely proportional to Bond Energy, so as Bond length gets shorter the bond energy gets larger

29. Bond Energies and Bond Lengths for Single Bonds
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Bond Energies and Bond Lengths for Single Bonds

30. Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Bond Length
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Visual Concept

31. Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Bond Energy
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Visual Concept

32. Noble Gases – Outer s and p orbitals are completely
The Octet Rule – Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an Octet of e- in its highest occupied energy level.
Noble Gases – Outer s and p orbitals are completely
filled with 8 electrons making them very
stable. (except He – has full 1s)
* Other main group atoms can effectively fill their
outermost s and p orbitals with e- by sharing e-
through covalent bonding by following the Octet
Rule.
*Fig. 9 p. 183
There are some exceptions to the Octet Rule (p. 183)

33. Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
The Octet Rule
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Visual Concept

34. - For the main group elements – don’t include the d block elements
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.
- For the main group elements – don’t
include the d block elements
- Must be in this order:
X X X X X X X X
Let’s Practice! P Si S Cl Xe

35. Section 2 Covalent Bonding and Molecular Compounds
Chapter 6

36. Lewis Structures
- formulas in which atomic symbols represent nuclei and inner-shell e- (F)
- Dot-pairs or Dashes between two atomic symbols represent e- pairs in covalent bonds
H - H or F : F
- Dots adjacent to only one atomic symbol represent unshared e-
:F : F: or :F - F:
- Structual Formula – indicates the kind, number, arrangement, and bonds but not the unshared pairs of atoms in a molecule.
F - F H-Cl
- Single Bond – a covalent bond produced by the sharing of one pair (two) electrons between two atoms
- Follow the steps in sample problem C p. 185 on how to write a Lewis Structure.

37. Lewis Structures Chapter 6
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Lewis Structures
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Visual Concept

38. Double Bond – covalent bond produced by the
Multiple Covalent Bonds - some elements, epecially Carbon, Nitrogen, and Oxygen can share more than one electron pair.
Double Bond – covalent bond produced by the
sharing of two pairs (4) electrons between
two atoms
Triple Bond – covalent bond produced by the
sharing of three pairs (6) electrons between
two atoms.
**The need for a multiple bond becomes obvious if
there are not enough valence e- to complete
octets by adding unshared pairs.
*Look at sample problem D p. 188
Resonance Structures – read p. 189

40. Comparing Single, Double, and Triple Bonds
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Comparing Single, Double, and Triple Bonds
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Visual Concept

41. Section 6-3 Ionic Bonding and Ionic Compounds
Objectives “I can” statements:
1) Compare a chemical formula for a molecular compound with one for an ionic compound. 
2) Discuss the arrangements of ions in crystals.
  3) Define lattice energy and explain its significance.
4) List and compare the distinctive properties of ionic and molecular compounds.

42. Section 6-3 Ionic Bonding and Ionic Compounds
Ionic Compound – composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal.
Most Ionic compounds exist as crystalline solids.
The chemical formula of an ionic compound represents the simplest ratio of the compound’s combined ions that gives electrical neutrality.
Na+ Cl-  NaCl
Molecular compounds are independent, neutral units with covalent bonds that can be isolated and examined
Formula Unit – the simplest collection of atoms from which an ionic compound’s formula can be established.
the ratio of ions in a formula unit depends on the charges of the ions combined. Ex: above with NaCl Ex: CaF2 (F- Ca2+)

43. Formation of Ionic Compounds
Ca F F  Ca F F- = (CaF2)
The Calcium atom has two valence electrons and each Fluorine atom has seven valence electrons.
Atoms of Calcium and other alkaline earth metals easily lose two electrons to form cations. 
- Atoms of Fluorine and other halogens easily gain one electron to form anions.
The transfer of two electrons from Ca to each F atom transforms each atom into an ION with noble-gas configuration.

45. Characteristics of Ionic Bonding
In an Ionic Crystal, ions minimize their Potential Energy (PE) by combining in an orderly arrangement known as CRYSTAL LATTICE.

46. NaCl and CsCl Crystal Lattices
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
NaCl and CsCl Crystal Lattices

48. Characteristics of Ion Bonding in a Crystal Lattice
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Characteristics of Ion Bonding in a Crystal Lattice
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Visual Concept

49. The distances between ions and their arrangement in a crystal represent a balance between attractive and repulsive forces
NaCl – in the crystal structure each Na is surrounded by 6 Cl and each Cl is surrounded by 6 Na for a 1:1 ratio. Fig. 15 p. 192.

50. Lattice Energy – the energy released when one mole of an ionic crystalline compound is formed from gaseous ions. This determines the IONIC BOND STRENGTH

51. Lattice Energy Chapter 6 Section 3 Ionic Bonding and Ionic Compounds
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Visual Concept

52. Comparision of Ionic and Molecular Compounds
The force that holds ions together in ionic compounds is a very strong ELECTROSTATIC attraction between positive and negative charges.
The forces between covalent bonds in molecules is much weaker than the forces in ionic bonding.
Ionic compounds generally have higher melting and boiling points than molecular compounds.
Ionic compounds are hard but brittle. A slight shift of one row of ions relative to another causes a large buildup of repulsive forces. If one layer is moved, the repulsive forces make the layers part completely and break (brittle). (Fig. 17 p. 193)

54. Melting and Boiling Points of Compounds
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Melting and Boiling Points of Compounds

55. Ionic Vs. Covalent Bonding
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Ionic Vs. Covalent Bonding

56. Many Ionic compounds are easily dissolved in water and become good electrical conductors. Don’t conduct electricity in the crystal state.

57. Comparing Ionic and Molecular Compounds
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Comparing Ionic and Molecular Compounds
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Visual Concept

58. Polyatomic Ions – a charged group of covalently bonded atoms have both molecular (covalent) and ionic characteristics
the charge of a polyatomic ion results from an excess of electrons (negative charge) or a shortage of electrons (positive charge).
examples on p. 194

59. Polyatomic Ions, continued
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Polyatomic Ions, continued
Some examples of Lewis structures of polyatomic ions are shown below.

60. Section 6-4 Metallic Bonding
Objectives or “I can” Statements:
1) Describe the electron-sea model of metallic bonding, and explain why metals are good electrical conductors.
2) Explain why metal surfaces are shiny.
3) Explain why metals are malleable and ductile but ionic-crystalline compounds are not.

61. Section 6-4 Metallic Bonding
Metallic Bonding – The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons.
Vacant orbitals in the atoms’ outer energy levels overlap.
This overlapping of orbitals allows the outer electrons of the atoms to roam freely throughout the entire metal.
The electrons are delocalized, which means that they do not belong to any one atom but move freely about the metal’s network of empty atomic orbitals.
These mobile electrons form a “sea of electrons” around the metal atoms which are packed together in a crystal lattice.

62. Metallic Bonding Chapter 6 Section 4 Metallic Bonding
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Visual Concept

63. Metallic Properties
High electrical and thermal conductivity results in metals due to the freedom of motion of electrons in a network.
Shiny appearance is the result of metals being able to absorb a wide range of light frequencies. Electrons are excited and fall back to lower E levels causing the shiny appearance.
Metallic bonding is the same in all directions through out the solid. One plane of atoms in a metal can slide past another without any resistance or breaking any bonds.
Allows special properties: Malleability – hammered
into thin sheets
Ductility – drawn or pulled to make wire

64. Properties of Metals: Surface Appearance
Section 4 Metallic Bonding
Chapter 6
Properties of Metals: Surface Appearance
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Visual Concept

65. Properties of Metals: Malleability and Ductility
Section 4 Metallic Bonding
Chapter 6
Properties of Metals: Malleability and Ductility
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Visual Concept

66. Properties of Metals: Electrical and Thermal Conductivity
Section 4 Metallic Bonding
Chapter 6
Properties of Metals: Electrical and Thermal Conductivity
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Visual Concept

67. Section 6-5 Intermolecular Forces Objectives or “I can” statements:
Describe dipole-dipole forces and hydrogen bonding, and their effects on properties such as boiling and melting points.
2) Distinguish between Intramolecular bonds that hold the inside of a molecule together and Intermolecular bonds that bond individual molecules together.

68. Section 6-5 Intermolecular Forces
Intermolecular Forces – forces of attraction between molecules
vary in strength but are generally weaker than the bonds that join the molecules (intramolecular forces). Add this
the higher the boiling point - the stronger the forces between particles.
Boiling points for ionic compounds and metals tend to be much higher than those for molecular substances: forces between molecules are weaker than those between metal atoms or ions.
Strongest Intermolecular forces exist between Polar Molecules

69. Comparing Ionic and Molecular Substances
Section 5 Molecular Geometry
Chapter 6
Comparing Ionic and Molecular Substances

70. Section 6-5 Intermolecular Forces
Dipole-Dipole Forces are the forces of attraction between polar molecules
- A dipole is created by equal but opposite charges that
are separated by a short distance.
- The direction of a dipole is from the dipole’s positive
pole to its negative pole.
- The negative region in one polar molecule attracts the
positive region in adjacent molecules. So the
molecules all attract each other from opposite sides.

71. Dipole-Dipole Forces Chapter 6 Section 5 Molecular Geometry
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Visual Concept

72. Section 6-5 Intermolecular Forces
Hydrogen Bonding – a strong type of Dipole-Dipole force. Water molecules have a high boiling point due to this strong bond.
Hydrogen bonding occurs between hydrogen atoms and very electronegative atoms which creates a POLAR bond. These elements are Nitrogen, Oxygen, and Fluorine. (N, O, F)

73. Visual Concepts
Chapter 6
Hydrogen Bonding

74. London Dispersion Forces
Section 5 Molecular Geometry
Chapter 6
Intermolecular Forces, continued
London Dispersion Forces
- In any atom or molecule—polar or nonpolar—the electrons are in continuous motion.
- As a result, at any instant the electron distribution may be uneven. A momentary uneven charge can create a positive pole at one end of an atom of molecule and a negative pole at the other.

75. London Dispersion Force
Section 5 Molecular Geometry
Chapter 6
London Dispersion Force
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Visual Concept

76. Intermolecular Forces, continued
Section 5 Molecular Geometry
Chapter 6
Intermolecular Forces, continued
London Dispersion Forces, continued
- This temporary dipole can then induce a dipole in an adjacent atom or molecule. The two are held together for an instant by the weak attraction between temporary dipoles.
- The intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles are called London dispersion forces.