1. Chapter 9 Chemical Bonding I: Lewis Theory
Chemistry: A Molecular Approach, 1st Ed. Nivaldo Tro
Chapter 9 Chemical Bonding I: Lewis Theory

2. Bonding Theories explain how and why atoms attach together
explain why some combinations of atoms are stable and others are not
why is water H2O, not HO or H3O
one of the simplest bonding theories was developed by G.N. Lewis and is called Lewis Theory
Lewis Theory emphasizes valence electrons to explain bonding
using Lewis Theory, we can draw models – called Lewis structures – that allow us to predict many properties of molecules
aka Electron Dot Structures
such as molecular shape, size, polarity
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3. TO LOWER THE POTENTIAL ENERGY
Why Do Atoms Bond?
TO LOWER THE POTENTIAL ENERGY
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4. Potential Energy Between Charged Particles
The repulsion between like-charged particles increases as the particles get closer together. To bring them closer requires the addition of more energy.
The attraction between opposite-charged particles increases as the particles get closer together. Bringing them closer lowers the potential energy of the system.
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5. Bonding
a chemical bond forms when the potential energy of the bonded atoms is less than the potential energy of the separate atoms
have to consider following interactions:
nucleus-to-nucleus repulsion
electron-to-electron repulsion
nucleus-to-electron attraction
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6. Types of Bonds Types of Atoms Type of Bond Bond Characteristic
metals to
nonmetals
Ionic
electrons
transferred
nonmetals to
Covalent
shared
metal to
metal
Metallic
pooled
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7. Types of Bonding

8. Ionic Bonds
when metals bond to nonmetals, some electrons from the metal atoms are transferred to the nonmetal atoms
metals have low ionization energy, relatively easy to remove an electron from
nonmetals have high electron affinities, relatively good to add electrons to
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9. Determining the Number of Valence Electrons in an Atom
Can be determined from the electron configuration.
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10. Determining the Number of Valence Electrons in an Atom
the column number on the Periodic Table will tell you how many valence electrons a main group atom has
Transition Elements all have 2 valence electrons; Why? 1A 2A 3A 4A 5A 6A 7A 8A Li Be B C N O F Ne
1 e-1
2 e-1
3 e-1
4 e-1
5 e-1
6 e-1
7 e-1
8 e-1
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11. Lewis Symbols of Atoms aka electron dot symbols
use symbol of element to represent nucleus and inner electrons
use dots around the symbol to represent valence electrons
pair first two electrons for the s orbital
put one electron on each open side for p electrons
then pair rest of the p electrons
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12. Lewis Symbols of Ions
Cations have Lewis symbols without valence electrons
Lost in the cation formation
Anions have Lewis symbols with 8 valence electrons
Electrons gained in the formation of the anion
Li• Li+1
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13. Lewis Theory
the basis of Lewis Theory is that there are certain electron arrangements in the atom that are more stable
octet rule
bonding occurs so atoms attain a more stable electron configuration
more stable = lower potential energy
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14. Stable Electron Arrangements And Ion Charge
Metals form cations by losing enough electrons to get the same electron configuration as the previous noble gas
Nonmetals form anions by gaining enough electrons to get the same electron configuration as the next noble gas
The noble gas electron configuration must be very stable
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15. Octet Rule
when atoms bond, they tend to gain, lose, or share electrons to result in 8 valence electrons
ns2np6
noble gas configuration
many exceptions
H, Li, Be, B attain an electron configuration like He
He = 2 valence electrons
Li loses its one valence electron
H shares or gains one electron
though it commonly loses its one electron to become H+
Be loses 2 electrons to become Be2+
though it commonly shares its two electrons in covalent bonds, resulting in 4 valence electrons
B loses 3 electrons to become B3+
though it commonly shares its three electrons in covalent bonds, resulting in 6 valence electrons
expanded octets for elements in Period 3 or below
using empty valence d orbitals
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16. Properties of Ionic Compounds
Melting an Ionic Solid
hard and brittle crystalline solids
all are solids at room temperature
melting points generally > 300C
the liquid state conducts electricity
the solid state does not conduct electricity
many are soluble in water
the solution conducts electricity well
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17. Conductivity of NaCl
in NaCl(s), the ions are stuck in position and not allowed to move to the charged rods
in NaCl(aq), the ions are separated and allowed to move to the charged rods
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18. Lewis Theory and Ionic Bonding
Lewis symbols can be used to represent the transfer of electrons from metal atom to nonmetal atom, resulting in ions that are attracted to each other and therefore bond
Li + +
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19. Predicting Ionic Formulas Using Lewis Symbols
electrons are transferred until the metal loses all its valence electrons and the nonmetal has an octet
numbers of atoms are adjusted so the electron transfer comes out even
2 Li +
Li2O
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20. Energetics of Ionic Bond Formation
the ionization energy of the metal is endothermic
Na(s) → Na+(g) + 1 e ─ DH° = +603 kJ/mol
the electron affinity of the nonmetal is exothermic
½Cl2(g) + 1 e ─ → Cl─(g) DH° = ─ 227 kJ/mol
generally, the ionization energy of the metal is larger than the electron affinity of the nonmetal, therefore the formation of the ionic compound should be endothermic
but the heat of formation of most ionic compounds is exothermic and generally large; Why?
Na(s) + ½Cl2(g) → NaCl(s) DH°f = -410 kJ/mol
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21. Ionic Bonds electrostatic attraction is nondirectional!!
no direct anion-cation pair
no ionic molecule
chemical formula is an empirical formula, simply giving the ratio of ions based on charge balance
ions arranged in a pattern called a crystal lattice
every cation surrounded by anions; and every anion surrounded by cations
maximizes attractions between + and - ions
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22. Lattice Energy
the lattice energy is the energy released when the solid crystal forms from separate ions in the gas state
always exothermic
hard to measure directly, but can be calculated from knowledge of other processes
lattice energy depends directly on size of charges and inversely on distance between ions
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23. Born-Haber Cycle
method for determining the lattice energy of an ionic substance by using other reactions
use Hess’s Law to add up heats of other processes
DH°f(salt) = DH°f(metal atoms, g) + DH°f(nonmetal atoms, g) + DH°f(cations, g) + DH°f(anions, g) + DH°f(crystal lattice)
DH°f(crystal lattice) = Lattice Energy
metal atoms (g)  cations (g), DH°f = ionization energy
don’t forget to add together all the ionization energies to get to the desired cation
M2+ = 1st IE + 2nd IE
nonmetal atoms (g)  anions (g), DH°f = electron affinity
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24. Born-Haber Cycle for NaCl
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25. Practice - Given the Information Below, Determine the Lattice Energy of MgCl2
Mg(s) ® Mg(g) DH1°f = kJ/mol
½ Cl2(g) ® Cl(g) DH2°f = kJ/mol
Mg(g) ® Mg+1(g) DH3°f = +738 kJ/mol
Mg+1(g) ® Mg+2(g) DH4°f = kJ/mol
Cl(g) ® Cl-1(g) DH5°f = -349 kJ/mol
Mg(s) + Cl2(g) ® MgCl2(s) DH6°f = kJ/mol
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26. Practice - Given the Information Below, Determine the Lattice Energy of MgCl2
Mg(s) ® Mg(g) DH1°f = kJ/mol
½ Cl2(g) ® Cl(g) DH2°f = kJ/mol
Mg(g) ® Mg+1(g) DH3°f = +738 kJ/mol
Mg+1(g) ® Mg+2(g) DH4°f = kJ/mol
Cl(g) ® Cl-1(g) DH5°f = -349 kJ/mol
Mg(s) + Cl2(g) ® MgCl2(s) DH6°f = kJ/mol
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27. Trends in Lattice Energy Ion Size
the force of attraction between charged particles is inversely proportional to the distance between them
larger ions mean the center of positive charge (nucleus of the cation) is farther away from negative charge (electrons of the anion)
larger ion = weaker attraction = smaller lattice energy
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28. Lattice Energy vs. Ion Size
Metal Chloride
Lattice Energy (kJ/mol)
LiCl
-834
NaCl
-787
KCl
-701
CsCl
-657
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29. Trends in Lattice Energy Ion Charge
-910 kJ/mol
the force of attraction between oppositely charged particles is directly proportional to the product of the charges
larger charge means the ions are more strongly attracted
larger charge = stronger attraction = larger lattice energy
of the two factors, ion charge generally more important
Lattice Energy =
-3414 kJ/mol
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30. Example 9.2 – Order the following ionic compounds in order of increasing magnitude of lattice energy. CaO, KBr, KCl, SrO
First examine the ion charges and order by product of the charges
Ca2+& O2-, K+ & Br─, K+ & Cl─, Sr2+ & O2─
(KBr, KCl) < (CaO, SrO)
Then examine the ion sizes of each group and order by radius; larger < smaller
(KBr, KCl) same cation, Br─ > Cl─ (same Group)
(CaO, SrO) same anion, Sr2+ > Ca2+ (same Group)
KBr < KCl < SrO < CaO
KBr < KCl < (CaO, SrO)
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31. Ionic Bonding Model vs. Reality
ionic compounds have high melting points and boiling points
MP generally > 300°C
all ionic compounds are solids at room temperature
because the attractions between ions are strong, breaking down the crystal requires a lot of energy
the stronger the attraction (larger the lattice energy), the higher the melting point
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32. Ionic Bonding Model vs. Reality
ionic solids are brittle and hard
the position of the ion in the crystal is critical to establishing maximum attractive forces – displacing the ions from their positions results in like charges close to each other and the repulsive forces take over + - + - + -
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33. Ionic Bonding Model vs. Reality
ionic compounds conduct electricity in the liquid state or when dissolved in water, but not in the solid state
to conduct electricity, a material must have charged particles that are able to flow through the material
in the ionic solid, the charged particles are locked in position and cannot move around to conduct
in the liquid state, or when dissolved in water, the ions have the ability to move through the structure and therefore conduct electricity
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34. Covalent Bonds
nonmetals have relatively high ionization energies, so it is difficult to remove electrons from them
when nonmetals bond together, it is better in terms of potential energy for the atoms to share valence electrons
potential energy lowest when the electrons are between the nuclei
shared electrons hold the atoms together by attracting nuclei of both atoms
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35. Single Covalent Bonds •• •• • • •• •• O H F O H F
two atoms share a pair of electrons
2 electrons
one atom may have more than one single bond H • O •• F •• • H O •• F •• F
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36. Double Covalent Bond •• • •• • ••
two atoms sharing two pairs of electrons
4 electrons O •• • O •• • O ••
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37. Triple Covalent Bond •• • •• • ••
two atoms sharing 3 pairs of electrons
6 electrons N •• • N •• • N ••
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38. Covalent Bonding Predictions from Lewis Theory
Lewis theory allows us to predict the formulas of molecules
Lewis theory predicts that some combinations should be stable, while others should not
because the stable combinations result in “octets”
Lewis theory predicts in covalent bonding that the attractions between atoms are directional
the shared electrons are most stable between the bonding atoms
resulting in molecules rather than an array
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39. Covalent Bonding Model vs. Reality
molecular compounds have low melting points and boiling points
MP generally < 300°C
molecular compounds are found in all 3 states at room temperature
melting and boiling involve breaking the attractions between the molecules, but not the bonds between the atoms
the covalent bonds are strong
the attractions between the molecules are generally weak
the polarity of the covalent bonds influences the strength of the intermolecular attractions
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40. Intermolecular Attractions vs. Bonding
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41. Ionic Bonding Model vs. Reality
some molecular solids are brittle and hard, but many are soft and waxy
the kind and strength of the intermolecular attractions varies based on many factors
the covalent bonds are not broken, however, the polarity of the bonds has influence on these attractive forces
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42. Ionic Bonding Model vs. Reality
molecular compounds do not conduct electricity in the liquid state
molecular acids conduct electricity when dissolved in water, but not in the solid state
in molecular solids, there are no charged particles around to allow the material to conduct
when dissolved in water, molecular acids are ionized, and have the ability to move through the structure and therefore conduct electricity
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43. Bond Polarity
covalent bonding between unlike atoms results in unequal sharing of the electrons
one atom pulls the electrons in the bond closer to its side
one end of the bond has larger electron density than the other
the result is a polar covalent bond
bond polarity
the end with the larger electron density gets a partial negative charge
the end that is electron deficient gets a partial positive charge
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44. HF
EN 2.1
EN 4.0
H F • d+ d-
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45. Electronegativity measure of the pull an atom has on bonding electrons
increases across period (left to right) and
decreases down group (top to bottom)
fluorine is the most electronegative element
francium is the least electronegative element
the larger the difference in electronegativity, the more polar the bond
negative end toward more electronegative atom
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46. Electronegativity Scale
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47. Electronegativity and Bond Polarity
If difference in electronegativity between bonded atoms is 0, the bond is pure covalent
equal sharing
If difference in electronegativity between bonded atoms is 0.1 to 0.4, the bond is nonpolar covalent
If difference in electronegativity between bonded atoms 0.5 to 1.9, the bond is polar covalent
If difference in electronegativity between bonded atoms larger than or equal to 2.0, the bond is ionic
0.4
2.0
4.0 4%
51%
Percent Ionic Character
Electronegativity Difference
“100%”

48. Bond Polarity ENCl = 3.0 ENH = 2.1 3.0 – 2.1 = 0.9 Polar Covalent
ENNa = 1.0
3.0 – 0.9 = 2.1
Ionic
ENCl = 3.0
= 0
Pure Covalent
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49. Tro, Chemistry: A Molecular Approach

50. Bond Dipole Moments
the dipole moment is a quantitative way of describing the polarity of a bond
a dipole is a material with positively and negatively charged ends
measured
dipole moment, m, is a measure of bond polarity
it is directly proportional to the size of the partial charges and directly proportional to the distance between them
m = (q)(r)
not Coulomb’s Law
measured in Debyes, D
the percent ionic character is the percentage of a bond’s measured dipole moment to what it would be if full ions
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51. Dipole Moments
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52. Water – a Polar Molecule
stream of water attracted to a charged glass rod
stream of hexane not attracted to a charged glass rod
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53. Example 9.3(c) - Determine whether an N-O bond is ionic, covalent, or polar covalent.
Determine the electronegativity of each element
N = 3.0; O = 3.5
Subtract the electronegativities, large minus small
(3.5) - (3.0) = 0.5
If the difference is 2.0 or larger, then the bond is ionic; otherwise it’s covalent
difference (0.5) is less than 2.0, therefore covalent
If the difference is 0.5 to 1.9, then the bond is polar covalent; otherwise it’s covalent
difference (0.5) is 0.5 to 1.9, therefore polar covalent
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54. Lewis Structures of Molecules
shows pattern of valence electron distribution in the molecule
useful for understanding the bonding in many compounds
allows us to predict shapes of molecules
allows us to predict properties of molecules and how they will interact together
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55. Lewis Structures B C N O F use common bonding patterns
C = 4 bonds & 0 lone pairs, N = 3 bonds & 1 lone pair, O= 2 bonds & 2 lone pairs, H and halogen = 1 bond, Be = 2 bonds & 0 lone pairs, B = 3 bonds & 0 lone pairs
often Lewis structures with line bonds have the lone pairs left off
their presence is assumed from common bonding patterns
structures which result in bonding patterns different from common have formal charges B C N O F
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56. Writing Lewis Structures of Molecules HNO3
Write skeletal structure
H always terminal
in oxyacid, H outside attached to O’s
make least electronegative atom central
N is central
Count valence electrons
sum the valence electrons for each atom
add 1 electron for each − charge
subtract 1 electron for each + charge
N = 5
H = 1
O3 = 3∙6 = 18
Total = 24 e-
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57. Writing Lewis Structures of Molecules HNO3
Attach central atom to the surrounding atoms with pairs of electrons and subtract from the total
Electrons
Start 24
Used 8
Left 16
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58. Writing Lewis Structures of Molecules HNO3
Complete octets, outside-in
H is already complete with 2
1 bond
and re-count electrons
N = 5
H = 1
O3 = 3∙6 = 18
Total = 24 e-
Electrons
Start 24
Used 8
Left 16
Electrons
Start 16
Used 16
Left 0
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59. Writing Lewis Structures of Molecules HNO3
If all octets complete, give extra electrons to central atom.
elements with d orbitals can have more than 8 electrons
Period 3 and below
If central atom does not have octet, bring in electrons from outside atoms to share
follow common bonding patterns if possible
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60. Practice - Lewis Structures
CO2
SeOF2
NO2-1
H3PO4
SO3-2
P2H4
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61. Practice - Lewis Structures
CO2
SeOF2
NO2-1
16 e-
H3PO4
SO3-2
P2H4
:O::C::O: :
32 e-
26 e-
26 e-
18 e-
14 e-
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62. FC = valence e- - nonbonding e- - ½ bonding e-
Formal Charge
during bonding, atoms may wind up with more or less electrons in order to fulfill octets - this results in atoms having a formal charge
FC = valence e- - nonbonding e- - ½ bonding e-
left O FC = ½ (4) = 0
S FC = ½ (6) = +1
right O FC = ½ (2) = -1
sum of all the formal charges in a molecule = 0
in an ion, total equals the charge • ••
O S O
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63. Writing Lewis Formulas of Molecules (cont’d)
Assign formal charges to the atoms
formal charge = valence e- - lone pair e- - ½ bonding e-
follow the common bonding patterns +1 -1
all 0
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64. Common Bonding PatternsF C B N O C + N + O + F + B - N - O - C - - F
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65. Practice - Assign Formal Charges
CO2
SeOF2
NO2-1
H3PO4
SO3-2
P2H4
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66. Practice - Assign Formal Charges-1
CO2
SeOF2
NO2-1
H3PO4
SO3-2
P2H4
P = +1
rest 0
all 0 -1 -1
Se = +1 -1
S = +1 -1 -1
all 0
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67. Resonance
when there is more than one Lewis structure for a molecule that differ only in the position of the electrons, they are called resonance structures
the actual molecule is a combination of the resonance forms – a resonance hybrid
it does not resonate between the two forms, though we often draw it that way
look for multiple bonds or lone pairs •• •
O S O
O S O
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68. Resonance
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69. Ozone Layer
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70. Rules of Resonance Structures
Resonance structures must have the same connectivity
only electron positions can change
Resonance structures must have the same number of electrons
Second row elements have a maximum of 8 electrons
bonding and nonbonding
third row can have expanded octet
Formal charges must total same
Better structures have fewer formal charges
Better structures have smaller formal charges
Better structures have − formal charge on more electronegative atom
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71. Drawing Resonance Structures
draw first Lewis structure that maximizes octets
assign formal charges
move electron pairs from atoms with (-) formal charge toward atoms with (+) formal charge
if (+) fc atom 2nd row, only move in electrons if you can move out electron pairs from multiple bond
if (+) fc atom 3rd row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet. -1 -1 +1 -1 +1
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72. Exceptions to the Octet Rule
expanded octets
elements with empty d orbitals can have more than 8 electrons
odd number electron species e.g., NO
will have 1 unpaired electron
free-radical
very reactive
incomplete octets
B, Al
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73. Drawing Resonance Structures
draw first Lewis structure that maximizes octets
assign formal charges
move electron pairs from atoms with (-) formal charge toward atoms with (+) formal charge
if (+) fc atom 2nd row, only move in electrons if you can move out electron pairs from multiple bond
if (+) fc atom 3rd row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet. -1 +2 -1
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74. Practice - Identify Structures with Better or Equal Resonance Forms and Draw Them-1
CO2
SeOF2
NO2-1
H3PO4
SO3-2
P2H4
P = +1
all 0 -1 -1
Se = +1 -1
S = +1 -1 -1
all 0
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75. Practice - Identify Structures with Better or Equal Resonance Forms and Draw Them
CO2
SeOF2
NO2-1
H3PO4
SO3-2
P2H4 -1
all 0
none +1 -1
all 0
S = 0
in all
res. forms +1 -1 -1
none
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76. Bond Energies
chemical reactions involve breaking bonds in reactant molecules and making new bond to create the products
the DH°reaction can be calculated by comparing the cost of breaking old bonds to the profit from making new bonds
the amount of energy it takes to break one mole of a bond in a compound is called the bond energy
in the gas state
homolytically – each atom gets ½ bonding electrons
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77. Trends in Bond Energies
the more electrons two atoms share, the stronger the covalent bond
C≡C (837 kJ) > C=C (611 kJ) > C−C (347 kJ)
C≡N (891 kJ) > C=N (615 kJ) > C−N (305 kJ)
the shorter the covalent bond, the stronger the bond
Br−F (237 kJ) > Br−Cl (218 kJ) > Br−Br (193 kJ)
bonds get weaker down the column
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78. Using Bond Energies to Estimate DH°rxn
the actual bond energy depends on the surrounding atoms and other factors
we often use average bond energies to estimate the DHrxn
works best when all reactants and products in gas state
bond breaking is endothermic, DH(breaking) = +
bond making is exothermic, DH(making) = −
DHrxn = ∑ (DH(bonds broken)) + ∑ (DH(bonds formed))
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80. Estimate the Enthalpy of the Following Reaction

81. Estimate the Enthalpy of the Following Reaction
H2(g) + O2(g) ® H2O2(g)
reaction involves breaking 1mol H-H and 1 mol O=O and making 2 mol H-O and 1 mol O-O
bonds broken (energy cost)
(+436 kJ) + (+498 kJ) = +934 kJ
bonds made (energy release)
2(464 kJ) + (142 kJ) = -1070
DHrxn = (+934 kJ) + ( kJ) = -136 kJ
(Appendix DH°f = kJ/mol)
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82. Bond Lengths
the distance between the nuclei of bonded atoms is called the bond length
because the actual bond length depends on the other atoms around the bond we often use the average bond length
averaged for similar bonds from many compounds
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83. Trends in Bond Lengths
the more electrons two atoms share, the shorter the covalent bond
C≡C (120 pm) < C=C (134 pm) < C−C (154 pm)
C≡N (116 pm) < C=N (128 pm) < C−N (147 pm)
decreases from left to right across period
C−C (154 pm) > C−N (147 pm) > C−O (143 pm)
increases down the column
F−F (144 pm) > Cl−Cl (198 pm) > Br−Br (228 pm)
in general, as bonds get longer, they also get weaker
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84. Bond Lengths
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85. Metallic Bonds
low ionization energy of metals allows them to lose electrons easily
the simplest theory of metallic bonding involves the metals atoms releasing their valence electrons to be shared by all to atoms/ions in the metal
an organization of metal cation islands in a sea of electrons
electrons delocalized throughout the metal structure
bonding results from attraction of cation for the delocalized electrons
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86. Metallic Bonding
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87. Metallic Bonding Model vs. Reality
metallic solids conduct electricity
because the free electrons are mobile, it allows the electrons to move through the metallic crystal and conduct electricity
as temperature increases, electrical conductivity decreases
heating causes the metal ions to vibrate faster, making it harder for electrons to make their way through the crystal
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88. Metallic Bonding Model vs. Reality
metallic solids conduct heat
the movement of the small, light electrons through the solid can transfer kinetic energy quicker than larger particles
metallic solids reflect light
the mobile electrons on the surface absorb the outside light and then emit it at the same frequency
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89. Metallic Bonding Model vs. Reality
metallic solids are malleable and ductile
because the free electrons are mobile, the direction of the attractive force between the metal cation and free electrons is adjustable
this allows the position of the metal cation islands to move around in the sea of electrons without breaking the attractions and the crystal structure
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90. Metallic Bonding Model vs. Reality
metals generally have high melting points and boiling points
all but Hg are solids at room temperature
the attractions of the metal cations for the free electrons is strong and hard to overcome
melting points generally increase to right across period
the charge on the metal cation increases across the period, causing stronger attractions
melting points generally decrease down column
the cations get larger down the column, resulting in a larger distance from the nucleus to the free electrons
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