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Chemical Bonding I: The Covalent Bond
Chapter 9
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Problems from pages 1-4, 5-26, 28, 30, 32, 34-36, 38, 40, 42, 44-46, 48, 50, 52, 54, 66
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Test on Chapters 7 and 8 April 6
EH Assignment Due April 4 Unit 11 Minimum score Sec 1 Size of Atoms and Ions 90 Sec 2 Ionization Energy & EA Sec 3 Metals, Nonmetals, Families Sec 4 Valence Electrons & Charge Sec 5 Acidity and Basicity 75 Test on Chapters 7 and 8 April 6
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Valence electrons are the outer shell electrons of an
atom. The valence electrons are the electrons that particpate in chemical bonding. Group # of valence e- e- configuration 1A 1 ns1 2A 2 ns2 3A 3 ns2np1 4A 4 ns2np2 5A 5 ns2np3 6A 6 ns2np4 7A 7 ns2np5 9.1
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The Ionic Bond - - - Li+ F Li + F 1s22s1 1s22s22p5 [He] 1s2 1s22s22p6
[Ne] Li Li+ + e- e- + F - F - Li+ + Li+ 9.2
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Covalent Bonding I Chemists use several different models to describe the covalent bond. These range from very simple to quite complex. The simplest is the Lewis Structure model which is discussed in detail in this chapter. It is very useful and successfully predicts the formulas and bonding in many molecules and polyatomic ions. Valence electrons are represented as dots. The octet rule is used. An atom tends to form bonds until it is surrounded by the same number of electrons as a noble gas.
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Covalent Bonding II In the Lewis structure model valence electrons are
presented as dots. The covalent bond is represented by a shared pair (bonding pair) of electrons. An electron pair that is not involved in bonding is called a lone pair or unshared pair or nonbonding pair. Bond Order: single, double, triple bonds. But there are many things this model cannot explain. In the next chapter the Valence Bond model and Molecular Orbital models are presented.
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Why should two atoms share electrons?
A covalent bond is a chemical bond in which two or more electrons are shared by two atoms. Why should two atoms share electrons? 7e- 7e- 8e- 8e- F F + F Lewis structure of F2 lone pairs F single covalent bond single covalent bond F 9.4
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Lewis structure of water
single covalent bonds 2e- 8e- 2e- H + O + H O H or Double bond – two atoms share two pairs of electrons 8e- 8e- 8e- double bonds O C or O C double bonds Triple bond – two atoms share three pairs of electrons 8e- triple bond N 8e- or N triple bond 9.4
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Lengths of Covalent Bonds
Bond Type Bond Length (pm) C-C 154 CC 133 CC 120 C-N 143 CN 138 CN 116 Bond Lengths Triple bond < Double Bond < Single Bond 9.4
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Last Test – April 26? EH Assignment Due April 15 Unit 12 Minimum score
Sec 1 Electronegativity and Bond Polarity 90 Sec 2 Lewis Dot Diagrams 88 Sec 3 Shapes of Ions and Molecules later Sec 4 Resonance and Formal Charge Sec 5 Review of Molecular Shapes Last Test – April 26?
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Electronegativity and Bond Polarity
Some bond are purely ionic - a complete transfer of electrons has occurred. NaCl Some bonds are pure covalent - the bonding electrons are shared equally. H2 or Cl2 In many cases the electrons are shared, but not equally. These bonds are called polar-covalent are considered to be partially ionic and partially covalent. HCl Electronegativity: the relative ability of a bonded atom to attract the shared electrons.
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Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms H F electron rich region electron poor region e- poor e- rich F H d+ d- 9.5
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Electron Affinity - measurable, Cl is highest
Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond. Electron Affinity - measurable, Cl is highest X (g) + e X-(g) Electronegativity - relative, F is highest 9.5
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Classification of bonds by difference in electronegativity
Bond Type Covalent 2 Ionic 0 < and <2 Polar Covalent Increasing difference in electronegativity Covalent share e- Polar Covalent partial transfer of e- Ionic transfer e- 9.5
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Classify the following bonds as ionic, polar covalent,
or covalent: The bond in CsCl; the bond in H2S; and the NN bond in H2NNH2. Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 Polar Covalent N – 3.0 N – 3.0 3.0 – 3.0 = 0 Covalent 9.5
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Writing Lewis Structures
Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center. Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge. Complete an octet for all atoms except H until all are complete or you run out of electrons. Form double and triple bonds on central atom if needed to complete octet. 9.6
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( ) - - Two possible skeletal structures of formaldehyde (CH2O) H C O
An atom’s formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure. formal charge on an atom in a Lewis structure = total number of valence electrons in the free atom - total number of nonbonding electrons - 1 2 total number of bonding electrons ( ) The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion. 9.7
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Formal Charge Formal charge: the charge an atom would have if the bonding electrons were shared equally. The formal charges must add up to equal the actual charge on the species. Atoms with a “normal” number of bonds have zero formal charge. Normal number of bonds - Cl, O, N, C, H
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Formal Charge and Lewis Structures
For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present. Lewis structures with large formal charges are less plausible than those with small formal charges. Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms. Which is the most likely Lewis structure for CH2O? H C O -1 +1 H C O 9.7
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For some molecules and ions no single Lewis
Resonance For some molecules and ions no single Lewis Structure can be drawn that agrees with experiment. In cases like this we can: Go on to a more sophisticated model of the covalent bond or 2. We can try to patch up the simple Lewis structure model to make it work.
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Resonance: the use of two or more Lewis
structures to represent a particular molecule or ion. Resonance structures are connected by a double-headed arrow. The true structure is a blend of the resonance structures. Resonance indicates that electron delocalization occurs in the molecule.
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What are the resonance structures of the carbonate (CO32-) ion?
A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure. O + - O + - What are the resonance structures of the carbonate (CO32-) ion? O C - O C - O C - 9.8
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Exceptions to the Octet Rule
The Incomplete Octet Be – 2e- 2H – 2x1e- 4e- BeH2 H Be B – 3e- 3F – 3x7e- 24e- 3 single bonds (3x2) = 6 9 lone pairs (9x2) = 18 Total = 24 F B BF3 9.9
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Exceptions to the Octet Rule
Odd-Electron Molecules N – 5e- O – 6e- 11e- NO N O The Expanded Octet (central atom with principal quantum number n > 2) S F S – 6e- 6F – 42e- 48e- 6 single bonds (6x2) = 12 18 lone pairs (18x2) = 36 Total = 48 SF6 9.9
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How the Model Explains the Properties
of Covalent Compounds The covalent bonding model proposes that electron sharing leads to strong, localized bonds. But most compounds with covalent bonds are gases, liquids, or low-melting solids. Ionic compound are high-melting solids. Why is this? Are covalent bonds weaker than ionic bonds? Within a molecule there are strong intramolecular bonds. Between molecules there are weak intermolecular forces. The latter are what determine melting and boiling points.
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Comparison of Ionic and Covalent Compounds
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Properties of the Period 3 Chlorides
Fig. 9.21
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For the test on chapter 9 study:
Test # questions 1-17 CHM Test #1 Spring 2001 Questions 1-8, 28, 29
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Single bond < Double bond < Triple bond
The enthalpy change required to break a particular bond in one mole of gaseous molecules is the bond energy. Bond Energy H2 (g) H (g) + DH0 = kJ Cl2 (g) Cl (g) + DH0 = kJ HCl (g) H (g) + Cl (g) DH0 = kJ O2 (g) O (g) + DH0 = kJ O N2 (g) N (g) + DH0 = kJ N Bond Energies Single bond < Double bond < Triple bond 9.10
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Average bond energy in polyatomic molecules
H2O (g) H (g) + OH (g) DH0 = 502 kJ OH (g) H (g) + O (g) DH0 = 427 kJ Average OH bond energy = 2 = 464 kJ 9.10
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Use bond energies to calculate the enthalpy change for:
H2 (g) + F2 (g) HF (g) DH0 = SBE(reactants) – SBE(products) Type of bonds broken Number of bonds broken Bond energy (kJ/mol) Energy change (kJ) H 1 436.4 F 1 156.9 Type of bonds formed Number of bonds formed Bond energy (kJ/mol) Energy change (kJ) H F 2 568.2 1136.4 DH0 = – 2 x = kJ 9.10
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