Chapter 1. Covalent Bonding and Shape of Molecules

Transcription

1 Learning objectives: hapter 1. ovalent Bonding and Shape of Molecules 1. Write the ground-state electron configuration. 2. Draw Lewis structure. 3. Use electronegativity to predict polarized and non-polarized bond. 4. Use VSEPR theory to predict the hybridization, bond angles, and geometry of an atom or molecule. 5. Recognize polar and non-polar molecules 6. Draw resonance structures by moving electron(s), and predict the order of stability for these resonance structures. 7. Recognize functional groups, such as alcohol, aldehyde, ketone, carboxylic acid, and amine. Sections to be covered: 1.1 Introduction # 1.2 Electron Structure of Atom 1.3 Lewis Model of Bonding * 1.4 Bond Angles and Shapes of Molecules * 1.5 Polar and Nonpolar Molecules 1.6 Resonance * 1.6 rbital verlap (Valence Bond) Model of ovalent Bonding * 1.7 Functional Groups * Sections that will be focused # Sections that will be skipped Recommended additional problems 1.17 to

2 1.2 Electron Structure of Atom A. Electron onfiguration of Atoms (i) Each atomic orbital can hold up to two electrons with their spin paired (opposite spin). energy spin paired (opposite spin) (ii) Fill electrons into atomic orbitals in order of increasing energy from lowest to highest. Know the energy level of basic atomic orbitals. (iii) For orbitals with equivalent energy, fill in one electron to each equivalent orbital before completely filling any one of these equivalent orbitals. energy equivalent orbitals energy equivalent orbitals 2

3 Examples: : e: Li: : N: : Na: P: B. Lewis Structures Know the number of valence electrons. Ne 1.3 Lewis Model of Bonding A. Formation of Ions (ctet rule, Anion and ation) Na Mg F 3

4 B. Formation of hemical Bond Ionic bond Nal ovalent bond 4. Electronegativity and hemical Bonds (i) Tendency of electronegativity on periodic table (ii) Important electronegativity : F: Li: : N: : 4

5 (iii) Electronegativity and chemical bonds Ionic bond Polar bond ovalent bond D. Drawing Lewis Structures of Molecules and Ions (i) Determine the number of valence electrons (ii) Know the number of legible (optimal) bonds and lone pair electrons for the commonly seen atoms: N F (l, Br, I) (iii) Determine the arrangement of atoms. Identify the enter Atom(s) (iv) Show the chemical bond as single line and non-bonding electrons as a pair of Lewis dots. 5

6 (v) Use multiple bonds when necessary double and triple bond (vi) Pay attention to the octet rule Examples: 2 N 3 4 l

7 E. Formal harge (i) Know the number of valence electrons for the commonly seen atoms. (ii) Know the number of legible (optimal) bonds and the corresponding charge for the commonly seen atoms. (iii) Arrange atoms according to the guidelines of section 1.3.D. Examples: 3 N 3 N

8 1.4 Bond Angles and Shapes of Molecules (i) Valence-shell electron-pair repulsion (VSEPR) model Know the total number of σ bond(s) and lone-pair electron(s) for the center atom(s) Examples: 4 N (ii) Difference of σ bond and π bond (iii) Lone-pair electrons will not be considered as part of the shape of molecule. (iv) The total number of σ bond(s) and lone-pair electron(s) for the center atom(s) dictate the arrangement of atoms surrounding the center atom(s) thus shape of molecules around the center atom(s). Total Number Description of Shape Predicted Bond Angles Possible Variation ybridization of enter Atom(s) 2 Linear 180 SP 3 Trigonal Planar 120 Linear-like (two pairs of lone-pair electrons) Bent (planar, one pair of lone-pair electrons) SP 2 4 Tetrahedral Bent (planar, two pairs of lone-pair electrons) Pyramidal (one pair of lone-pair electrons) SP 3 8

9 1.5 Polar and Nonpolar Molecules Examples: (i) 2 (ii) 2 (iii) N 3 (iv) F 4 (v) 3 l 9

10 1.6 Resonance A. Theory and Representation of Resonance (i) Molecules could have more than one structure. (ii) π electrons and lone-pair electrons are potentially delocalizable (movable) but not the σ electrons. (iii) Indication of electron movement using curved arrow. Electron-pushing or arrow-pushing B. Rules for Writing acceptable Resonance Structures - Same number of valence electrons - Same position (connectivity) of atoms - bey rule of covalent bonding: No more than 2 valence electrons for 1 st row elements (, e). No more than 8 valence electrons (octet rule) for 2 nd row elements. 3 rd row elements (P, S) may have more than 8 valence electrons. 10

11 Examples: N 3 S 3 3 N 2 11

12 . Effect of Resonance (Supplementary Material) (i) More resonance structure, better stability (ii) Prediction of electron density distribution based on the stability among resonance structures (in order of significance): - Neutral molecule is generally more stable than charged molecule. - Molecule having all the 2 nd atoms meet the octet rule is more stable than the one that does not. - Negative charge locates on the atom with high electronegativity is preferred. - Wider separation of positive and negative charges is preferred. (iii) hemical properties of molecules can be predicted by resonance effect Examples: (i) Why ethanoic acid (acetic acid), 3 2 is a stronger acid than methanol, 3? (ii) Which atom in the structure below has the highest (or lowest) electron density? 2 (iii) utline the possible resonance structures and predict the order of stability. 2 12

13 1.6 rbital verlap (Valence Bond) Model of ovalent Bonding A. Shapes of Atomic rbitals 1S 2S 2P (2P x, 2P y, 2P z ) B. Formation of a ovalent Bond by the verlap of Atomic rbitals Types and characteristics of σ bond and π bond. ybridization of Atomic rbitals (Also see section 1.4) ybridization reduces the electron repulsion and increases the stability. D. SP 3 ybrid rbitals 4 N

14 E. SP 2 ybrid rbitals BF 3 2 F: SP ybrid rbitals

15 1.7 Functional Groups - -N