1. General Bonding Concepts
Chapter 8 Sections 1-12 only

2. Objectives
Define 3 different types of bonds based on electron distribution
Use Coulomb's law equation to determine attraction/repulsion

3. Chemical Bonds What is a bond? Why do bonds form? What is bond energy?
The forces that cause atoms to act in unison
Why do bonds form?
to use or achieve the lowest possible energy
What is bond energy?
the amount of energy required to BREAK the bond

4. Types of Bonds Ionic: Electrons are transferred
Polar Covalent: Electrons are shared unequally
Covalent: Electrons are shared equally

5. Ionic Bonding
What types of substances react to form ionic bonds? Answer: Metals and non metals Energy of interaction between a pair of ions can be calculated using COULOMB'S LAW. The formula:

6. Coulomb’s Law Equation
Constant:
Joules= energy unit, nm = distance unit
Charge on Ions
Distance between atoms (in nm)
In Joules

7. Using Coulomb's Law
In NaCl the distance between ions is nm, What is the energy of this bond interaction? What equation? Set it up: Answer: E = -8.37x10-19 J

8. Coulomb's Law
A Negative Coulomb's Law answer indicates an attractive force
A Positive Coulomb's Law answer indicates a repulsive force

9. Coulomb’s Law Practice

10. Objectives Discuss relationship between bond energy and bond length
Discuss how electronegativity impacts bonding
Determine if a bond will be polar, if so show dipole moment
Define bond, bond energy, bond length, dipole moment
Explain the three forces at work when atoms bond

11. Relate Bond Energy and Bond Length
Finish this statement:
as bond length decreases bond energy…
INCREASES
Defend using Coulomb's Law:

12. What about two same type atoms?
Example H:H
A bond will form if the energy of the aggregate is lower than that of the separated ions or atoms (p. 331)
Picture from:
chemkb.cs.iupui.edu

13. Bonding between same type ions
These are covalent bonds.
Bond Length is defined as the distance where the lowest energy is achieved.
Several forces are involved to establish bond length

14. The Forces at Work When same type atoms interact you get a covalent bond. The following forces must be balanced for a covalent bond to occur:
What is this trying to do?
Type of interaction:
Proton vs Proton
Proton vs Electron
Electron vs Electron
Repel
Attract

15. Key points of Covalent Bonding
Shared Electrons
Electrons shared in area between nuclei
Bonds created to increase stability

16. Key points of Covalent Bonding
Unequal sharing of electrons
Based on electronegativity difference
Electronegativity Differences and bond type
0-1 "covalent"
1-2 "polar covalent"
>2 "ionic“
Actually more like a sliding scale
Most Covalent Most Ionic

17. Consequences of Unequal Sharing
Dipole Moment: partial charges caused by the unequal sharing of electrons
Polar bond does not always equal polar molecule
Dipoles can cancel due to molecular structure
Two ways to show dipoles:
H F
H F d+ d-

18. Assignment
Page 383 # 24, 26, 30, 32

19. Show the bond polarity for the following bonded atoms:
C-O Se-S P-H
Cl-I H-Cl Br-Te
Br-Br O-P Si-S

20. Objectives: Describe what it means for an atom/ion to be stable
Use electron configurations to predict formulas of ionic compounds
Describe the relationship between parent atom size and ion size for cations and anions
Use a periodic trend to determine relative ion size in a group
Define an isoelectronic ion group
Discuss the trend for ions size in an isoelectronic group

21. Stable Atoms
Quantum mechanical model has helped to show that atoms in a stable compound have achieved noble gas configuration (full energy level) by either sharing electrons or by forming ions by either the loss or gain of electrons.

22. Stable covalent compounds are achieved by atoms sharing electrons so that both complete their valence shells to attain noble gas configuration
Stable ionic compounds are formed when a non-metal removes the valence electrons from the metal so that it fills its valence shell to achieve noble gas configuration, the metal reverts back to the last noble gas configuration.

23. Predicting Ion Formation
You can use electron configurations to predict ion formation
Ca: [Ar]4s2
Ca wants to lose 2 electrons and achieve the configuration of Argon
O: [He] 2s22p4
O wants to gain 2 electrons and achieve the configuration of Neon
From this we predict:
The calcium ion The oxide ion The compound calcium oxide
Ca2+ O2- CaO

24. Ion Prediction and Compound Formation
What happens when aluminum and oxygen react?
Look at electron configurations:
Al = [Ne] 3s23p1 O = [He] 2s22p4
What does each atom "want"?
Al "wants" to lose 3 electrons,
O "wants" to gain 2 electrons
What formula would you predict? Why?
Al2O3
all atoms are "satisfied"

25. Ion Size
Important when determining stability, structure, and properties of ionic solids and aqueous ions.
Impossible to precisely define ion size
There are several ways to look at trends for ion size

26. Ion Size and The Parent Atom
Cations are smaller than their parent atom. Why?
loss of electrons results in the loss of an entire energy level which decreases the size of the entity
Anions are larger than their parent atom.
Why?
gain of electrons results in a fuller outer energy level which increases the size of the entity

27. Ion Size and Periodic Groups
Ion size increases down a group.
Why?
as you move down a group, there are more filled energy levels= larger ion
There is not a trend for ion size across a period like there is for other types of periodic trends.
as you move across a period, there is a change over from metals (cations) to non metals (anions) so in the middle of a period, it will change from smaller ions to large ones

28. Ion Size and Isoelectronic Ions
Define Isoelectronic:
ions containing the same number of electrons
Example:
O-2, F-1, Na+1, Mg+2, Al+3
What is the electron configuration of each ?
They all have the configuration of NEON!
10 total electrons are in each ion!
This is an isoelectronic set of ions

29. Ion Size and Isoelectronic Ions ... continued
The trend for isoelectronic ions is that ion size decreases as atomic number (nuclear charge, number of protons, z) increases.
Why??
Isoelectronic ions by definition have the same number of electrons, therefore, the greater the number of protons (atomic number) the greater the positive force drawing those electrons towards the nucleus (making the ion smaller by drawing the electrons closer)

30. Ion Size and Isoelectronic Ions... Practice
Arrange the following in order of decreasing ion size
Br-1, Rb+1, Se-2, Sr+2
Se-2, Br-1, Rb+1, Sr+2
Choose the smallest ion from the following sets
Li+1, Na+1, K+1, Rb+1
Li+1
Ba+2, Cs+1, I-1, Te-2
Ba+2

31. Assignment:
Page 383 # 35, 37, 39, 41

32. Answers:
#35 a. Sc+3 , b. Te-2 , c. Ce+4 and Ti+4, d. Ba+2 #37. La+3 , Ba+2 , Cs+1 , I-1 , Te-2 #39 a. Cu > Cu+1 > Cu+2 b. Pt+2 >Pd+2 >Ni+2 c. O-2 , O-1 , O d. La+3 , Eu+3 , Gd+3 , Yb+3 e . Te-2 , I -1 , Cs+1 , Ba+2 , La+3 #41 a. Al2S3 – aluminum sulfide b. K3N – potassium nitride c. MgCl2 – magnesium chloride d. CsBr – cesium bromide

33. Objectives: Define lattice energy
Determine relative lattice energy among a group of compounds
Use lattice energy to calculate reaction energy (5 step)

34. Na+1 (g) + Cl-1 (g) --> NaCl (s)
Lattice Energy
Define Lattice Energy:
the change in energy that takes place when separated gaseous ions are packed together to form an ionic solid
Na+1 (g) + Cl-1 (g) --> NaCl (s)
change in energy from here
to here
is lattice energy !

35. The Textbooks's Perspective on Energy
Viewed from the system's point of view
Negative energy if the process is exothermic
Positive energy if the process is endothermic
This perspective gives lattice energy a negative value!

36. Lattice Energy Calculation
Lattice Energy is Calculated by a modified form of Coulomb's Law.
Where k is a constant that depends on the structure of the solid and the electron configurations of the ions
Note: you will NOT be expected to calculate lattice energy, just to know how it is calculated and to use lattice energy to calculate total reaction energy

37. Estimating Relative Lattice Energy
Lattice energy increases as ion charge (Q1 and Q2) increases
Lattice energy decreases as the distance between ions (r) increases

38. Practice:
Which compound in the following pairs has the most exothermic lattice energy? Why?
NaCl or KCl
Mg(OH)2 or MgO
LiF or LiCl
Fe(OH)2 or Fe(OH)3
NaCl or Na2O
MgO or BaS
NaCl : smaller r
MgO : larger Q
LiF : smaller r
Fe(OH)3 : larger Q
Na2O : Larger Q
MgO : smaller r

39. Ionic Solid Formation Energy Calculation
Li (s) + 1/2 F2 (g) --> LiF (s) Note: must have balanced equation
Use lattice energy to determine the energy change experienced in this reaction.
The calculation is broken into steps based on states of matter
Sublimation
Ionization
Dissociation
Ion formation
Solid formation (lattice energy**)
** remember to use lattice energy you need separate, gaseous ions. The other 4 steps just get you to this point.

40. An Example Step Process Energy change Sublimation Li (s) --> Li (g)
161 kJ/mol
Ionization
Li (g) --> Li e-
520 kJ/mol
Dissociation
1/2 F2 (g) --> F (g)
154 kJ/mol
Ion Formation
F (g) + e- --> F-1 (g)
-328 kJ/mol
Solid Formation
Li+1(g) + F-1 (g) --> LiF (s)
-1047 kJ/mol
Total Energy
Li (s) + 1/2 F2 (g) --> LiF (s)
-617 kJ
the negative sign tells that energy is being released (exothermic)

41. Assignment:
Page # 44, 45, 46, 49

42. Answers:

46. A better definition of ionic compounds:
Any compound that conducts electricity when melted is ionic
Avoids confusion when polyatomic ions, which are often held together covalently are part of the compound

47. Use of Models Models attempt to explain how nature operates
Models are human inventions: Models do not equal reality
Models can be wrong. They are oversimplifications. Exceptions deal with items that do not meet standards due to oversimplifications.
Need to understand strengths and weaknesses
When models are wrong, we can learn

48. Using Bond Energy to Calculate Reaction Energy
Bond Energy depends on the environment
Example:
C-H bond in HCCl3 = 380 kJ/mol
C-H bond in C2H6 = 410 kJ/mol
Average Bond Energy is still useful

49. Objectives: Discuss a "bond" as a model
Explain why models are used in science
Discuss 5 properties of models
Define single, double and triple bonds- # of e- shared, bond energy & length
Use bond energies to calculate ΔH
Explain the relationship between number of shared electrons and bond strength and/or bond length
Write the expression for ΔH when using bond energy to calculate the enthalpy of the reaction

50. Ways to share electrons
Single Bonds single pair of electrons shared
C-C - two total shared electrons
Double Bonds two pairs of electrons shared
C=C four total shared electrons
Triple Bonds - three pairs of electrons shared
C=C - six total shared electrons

51. Using bond energy to calculate Enthalpy
H2 (g) + F2 (g) --> 2 HF (g)
This reaction requires us to break a H-H bond and a F-F bond on the reactant side of the reaction It also requires that we make two H-F bonds on the product side of the reaction.
Breaking bonds is an endothermic process- positive numbers
Making bonds is an exothermic process- negative numbers
Average bond energy Table 8.4 is on page 351.

52. How it is done . . .
Enthalpy change for a reaction is calculating by the addition of the energy required to break old bonds and the energy released by the formation of new bonds.
ΔH = ΣD (bonds broken) - ΣD (bonds formed)
Note: you are actually adding the energy from "bonds formed", however, since these are releasing energy the sign on the numbers is negative leading to the negative sign in the equation.

53. Back to the problem . . . H2 (g) + F2 (g) --> 2 HF (g)
Breaking one mol 432 kJ/mol = 432 kJ
Breaking one mol 154 kJ/mol = 154 kJ
Making two moles 565 kJ/mol = 1130 kJ
Bonds broken total= 586 kJ
Bonds formed total= 1130 kJ
= -544 kJ
Average bond energy Table 8.4 is on page 351.

54. Try this one on your own . . . carbon is central atom
CH4 + 2 Cl2 + 2 F2 --> CF2Cl2 + 2HF + 2HCl
Answer: kJ

55. Assignment:
Page 384 # 53,55,57,59,61
Answers

56. Answers
Page 384 # 53,55,57,59,61

57. Objectives Explain the Localized Energy Model Define lone pair
Define bonding pair
Define Lewis structure
Define duet rule
List elements that obey the duet rule
Define Octet rule
List elements that obey the octet rule

58. Objectives Draw Lewis structures for atoms
Draw Lewis structures for molecules
Explain how there are exceptions to the octet rule
List some elements that are exceptions to the octet rule Draw Lewis structures for compounds where elements exceed the octet rule

59. The Localized Electron Model
**Assumes that a molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms

60. The LE Model Continued
Electrons are assumed to be localized on an atom or in the space between two atoms.
Lone Pair : Localized on one atom
Bonding Pair: Found in space between two atoms

61. Three Parts of LE Model
Description of valence electron arrangement with Lewis Structures
Prediction of Molecular Geometry using VSEPR Model
Description of the type of atomic orbitals used by atoms to share electrons.
For this unit we are only going to deal with the first part, the second and third will be worked with next unit.

62. Lewis Structures
Shows how valence electrons are arranged among atoms in a molecule.
Most important requirement for a stable compound is the achievement of noble gas configuration.

63. Two Basic Rules
Duet Rule: a stable molecule is formed when two electrons are shared
Octet Rule: a stable molecule is formed when electrons are shared so that each atom is surrounded by eight electrons

64. Steps for Writing Lewis Structures
Sum the VALENCE electrons for all atoms in the molecule, it does not matter how many come from each, just the total
Use a pair of electrons to form a bond between each pair of bond atoms
Arrange remaining atoms to satisfy the duet rule for hydrogen and the octet rule for second row elements.

65. An Example:
Water: H2O
1. Sum valence electrons: Total= 8 valence electrons
H= 1 valence electron
O= 6 valence electrons
2. Use a pair to form bond between atoms
3. Arrange remaining to satisfy duet/octet rule

66. Another Example: Carbon Dioxide: CO2 Sum valence electrons
Total: 16 valence electrons
C= 4 valence electrons
O= 6 valence electrons
2. Use pairs to bond atoms
3. Arrange remaining to meet duet/octet

67. Assignment
Page 385 # 67 and 68
Answers

68. Answers
Page 385 # 67 and 68

69. Violations of the Octet Rule
Be, B are often electron deficient
Third row and beyond elements are able to exceed the octet rule by placing extra electrons in their "d" orbitals

70. Examples:
BF3
SF6
PCl5

71. Assignment
Page , 71,72
Answers

72. Answers
Page , 71,72

73. Objectives Define Resonance Define Resonance structure
Explain why odd electron molecules cannot be shown using the Localized Energy Model
Define formal charge
Explain the “problem” with oxidation state assignments
Explain how formal charge is different from oxidation state
Use formal charge to determine the most probable Lewis structure for a molecule

74. Resonance
Occurs when there is more than one equivalent and valid Lewis structure for a molecule.
Is necessary to compensate for the incorrect assumption that electrons are localized
In reality, electrons are delocalized- meaning that they move around the entire molecule
LE model is still useful, so the exception of resonance is added to accommodate certain molecules rather than tossing the whole model

75. An example:
NO3-1

76. Another Example
NO2-1

77. Assignment
Page 385 # 73,74,75
Answers

78. Answers
Page 385 # 73,74,75

79. Odd - Electron Molecules
Only a few molecules form containing odd numbers of electrons
Examples: NO and NO2
The Localized Electron Model only works with PAIRS of electrons so this model is not able to accomodate odd electron molecules

80. Formal Charge
Formal charge works with nonequivalent Lewis structures often found in molecules and polyatomic ions that have atoms that can exceed the octet rule.
Charges are utilized to determine the most appropriate structure(s)

81. Charge Determination Oxidation states:
Both shared electrons count for the more electronegative atom
Results in exaggerated charge estimates
Not good for determining proper Lewis structures

82. Charge Determination Formal Charge:
The difference between the number of valence electrons on the free atom and the number of valence electrons assigned to the atom in the molecule
When calculating formal charge:
Lone pair electrons belong to the atom on which they are located.
Shared pair electrons are split equally among sharing atoms.

83. Rules for working with Formal Charge:
Atoms in molecules want to have formal charges as close to zero as possible.
Negative formal charges should occur on the more electronegative atom.

84. An Example:
SO4-2

85. Another Example
XeO3
Hint: there are 8 possible Lewis structures

86. Assignment
Page 386 # 81 and 82
Answers

87. Answers
Page 386 # 81 and 82

88. End of Unit 1
Chapter 8 Sections 1-12.