Chapter 10 Basic Concepts of Chemical Bonding

Transcription

1 Chapter 10 Basic Concepts of Chemical Bonding 1 I. Covalent Bonds and Lewis Bonding Theory A. Lewis Symbols, the Octet Rule and Covalent Bonding B. Other Types of Covalent Bonds 1. Multiple Covalent Bonds 2. Coordinate Covalent Bonds C. Sharing of Electrons in Covalent Bonds 1. Electronegativity 2. Bond Polarity II. Molecular Compounds and Lewis Structures in Detail A. Drawing Lewis Structures 1. Formal Charge 2. Exceptions to the Octet Rule 3. Resonance Structures III. Molecular Shapes and Polarity A. VSEPR Theory B. Polarity of Molecules

2 I. Covalent Bonds and Lewis Bonding Theory Types of Compounds 2 1. Ionic Compounds Ionic Bonds - transfer of electrons between atoms metal - non-metal 2. Molecular Compounds Covalent Bonds - non-metal - sharing of electrons between atoms non-metal CH 3 COOH

3 3 A. Lewis Symbols, the Octet Rule and Covalent Bonding How do atoms interact? Which electrons can be shared? -they share electrons -valence electrons (outer s & p e - ) Lewis Bonding Theory Observations atoms react until they obtain a noble gas core Octet Rule most atoms share electrons until they are surrounded by 8 valence electrons Duet Rule Hydrogen and Helium share electrons until they are surrounded by 2 electrons Lewis Symbol chemical symbol for element + a dot for each valence electron

4 Main group elements and valence e - Group Element E.C. 4 Lewis Symbol 1 H 1s 1 H 13 B 1s 2 2s 2 2p 1 B 15 P [Ne]3s 2 3p 3 P 17 Br [Ar] 3d 10 4s 2 4p 5 Br Covalent Bonding, Lewis Symbols and Lewis Structures Covalent Bond - involves the sharing of electrons between two atoms Lewis Structure - structure obtained by combining Lewis symbols to create covalent bonds

5 B. Other Types of Covalent Bonds 5 1. Coordinate Covalent Bonds - covalent bond in which both e - of the shared pair are donated by one atom Example

6 2. Multiple Covalent Bonds - covalent bond formed by the sharing of more than 2 electrons (in multiples of 2) 6 Type of Bond Single e - shared Example Length Strength Double Triple Examples of Lewis Structures with Multiple Bonds

7 C. Sharing of Electrons in Covalent Bonds 7 1. Electronegativity - Are the electrons shared equally between the two N atoms? - Are the electrons shared equally between the C and the O atoms? How can we predict the equality of sharing? Which atoms want electrons the most?

8 2. Bond Polarity - measure of how well atoms share e - in a bond - calculated using electronegativity Determining Bond Polarity e.n. = e.n. of atom w/ largest e.n. - e.n. of atom with smallest e.n. 8 Approximations 1. Δe.n. below ~ Δe.n. between ~0.5 and~2 3. Δe.n. above ~2

9 Examples 9

10 II. Molecular Compounds and Lewis Structures in Detail 10 Requirements and Trends for L.S. 1. L.S. must include all valence e- - + ion subtract e- - - ion add e- 2. Usually e- are paired 3. Octet rule is usually followed (many exceptions) 4. Trends

11 5. Skeletal Structure - arrangement of atoms in molecule Central atom bonded 2 or more atoms Terminal atom bonded to 1 atom - H-atoms are always terminal - Carbon atoms are almost always central atoms - central atoms generally have the lowest electronegativity - More compact and symmetrical the better examples CO CH 3 CH 2 OH a497-b19d59590fef-jpeg jpg

12 A. Writing Lewis Structures 12 Plan of Attack 1. Sum all of the valence electrons from all of the atoms 2. Draw the skeleton structure of the molecule using one pair of electrons (a single bond) for each connecting bond. (If there is a central atom it is usually written first in the formula and usually has the lowest e.n.) 3. Determine the amount of electrons used and the number of electrons left over

13 4. Distribute the remaining electrons to achieve a noble gas configuration (octet) on each atom: 1) How many are needed? 2) How many are available? 13 -If needed = available distribute as lone pairs -If needed > available make one additional bond for every 2 electrons short and distribute remaining electrons as lone pairs Qualitative Approach

14 1. Formal Charge (a quantitative approach) 14 Formal Charge The hypothetical charge on an atom in a molecule assuming equal sharing of the bonding electrons F.C. # valencee - - lone e 1 - (shared e ) 2 Rules 1. formal charges of all atoms in molecule must add up to the charge of the molecule or ion 2. For non-equivalent Lewis Structures, in the best structure

15 Example Which is the best Structure? : O C N : or : O C N : or : O C N : A. B. C. Formal Charges A. B. O = 6 [4 + ½(4)] = 0 C = 4 [0 + ½(8)] = 0 N = 5 [4 + ½(4)] = -1 C. O = 6 [6 + ½(2)] = -1 C = 4 [0 + ½(8)] = 0 N = 5 [2 + ½(6)] = 0

16 Lewis Structure Example 16

17 2. Exceptions to the Octet Rule a. Odd # s of electrons b. Less than an octet c. More than an octet 17 a. Odd # s of Electrons

18 b. Less than an Octet 18 c. More than an Octet (Most common) Experimental central atoms with more than octet Explanation Period 1: Period 2: Period 3:

19 Example 19

20 3. Resonance Structures What happens if there are multiple equally likely Lewis Structures possible? 20