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Chapter 6

CHM 138 BASIC CHEMISTRY. Chapter 6. CHEMICAL BONDS. NOR AKMALAZURA JANI. Group. e - configuration. # of valence e -. ns 1. 1. 1A. 2A. ns 2. 2. 3A. ns 2 np 1. 3. 4A. ns 2 np 2. 4. 5A. ns 2 np 3. 5. 6A. ns 2 np 4. 6. 7A. ns 2 np 5. 7.

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Chapter 6

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  1. CHM 138 BASIC CHEMISTRY Chapter 6 CHEMICAL BONDS NOR AKMALAZURA JANI

  2. Group e- configuration # of valence e- ns1 1 1A 2A ns2 2 3A ns2np1 3 4A ns2np2 4 5A ns2np3 5 6A ns2np4 6 7A ns2np5 7 Valence electrons are the outer shell electrons of an atom. The valence electrons are the electrons that participate in chemical bonding.

  3. Lewis Dot Symbols for the Representative Elements & Noble Gases • Lewis dot symbol consists of the symbol of an element and one dot for each valence electron in an atom of the element.

  4. - + Li+ Li F F The Ionic Bond • Ionic bond: the electrostatic force that holds ions together in an ionic compound. • Atoms of the elements with low ionization energies tend to form cation – alkali metals and alkaline earth metals • Atoms of the elements with high electron affinities tend to form anion – halogens and oxygen (LiF) 1s22s1 1s22s22p5 1s2 1s22s22p6

  5. - - - Li Li+ + e- e- + Li+ Li+ + F F F F

  6. + 8e- 7e- 7e- 8e- F F F F F F F F lonepairs lonepairs single covalent bond single covalent bond lonepairs lonepairs A covalent bond is a chemical bond in which two or more electrons are shared by two atoms. Why should two atoms share electrons? Lewis structure of F2

  7. A Lewis structure: - a representation of covalent bonding in which shared electron pairs are shown either as lines or as pairs of dots between two atoms, and lone pairs are shown as pairs of dots an individual atoms. • Only valence electrons are shown. • The formation of the molecules illustrates the octet rule. - Octet rule: An atom other than hydrogen tends to form bonds until it is surrounded by eight valence electrons.

  8. single covalent bonds H H H H or H H O 2e- 2e- O 8e- O C O C O O double bonds 8e- 8e- 8e- O N N N N triple bond 8e- 8e- Lewis structure of water + + Double bond – two atoms share two pairs of electrons or double bonds Triple bond – two atoms share three pairs of electrons or triple bond

  9. Lengths of Covalent Bonds Bond Lengths Triple bond < Double Bond < Single Bond

  10. COMPARISON OF GENERAL PROPERTIES OF IONIC COMPOUND AND COVALENT COMPOUND

  11. Writing Lewis Structures • Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge. • Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center. • Complete an octet for all atoms except hydrogen • If structure contains too many electrons, form double and triple bonds on central atom as needed.

  12. F N F F Write the Lewis structure of nitrogen trifluoride (NF3). Step 1 – N is less electronegative than F, put N in center Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5) 5 + (3 x 7) = 26 valence electrons Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms. Step 4 - Check, are # of e- in structure equal to number of valence e- ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

  13. O C O 2 single bonds (2x2) = 4 1 double bond = 4 8 lone pairs (8x2) = 16 O Total = 24 Write the Lewis structure of the carbonate ion (CO32-). Step 1 – C is less electronegative than O, put C in center • Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4) • -2 charge – 2e- 4 + (3 x 6) + 2 = 24 valence electrons Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms. Step 4 - Check, are # of e- in structure equal to number of valence e- ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons Step 5 - Too many electrons, form double bond and re-check # of e-

  14. Write the Lewis structure for: i) NO -2 ii) CS2 iii) SO3

  15. formal charge on an atom in a Lewis structure total number of valence electrons in the free atom - = FORMAL CHARGE 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. total number electron assigned to atom The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion.

  16. H C O H H C O H Examples: formal charge on C = -1 formal charge on O = + 1 formal charge on C = 0 formal charge on O = 0

  17. H C O H H C O H -1 +1 0 0 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?

  18. - - + + O O O O O O O O O C C C O O O - - - - O O O - - Resonance Structure 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. What are the resonance structures of the carbonate (CO32-) ion?

  19. Be – 2e- 2H – 2x1e- H Be H 4e- B – 3e- 3 single bonds (3x2) = 6 3F – 3x7e- F B F 24e- Total = 24 9 lone pairs (9x2) = 18 F Exceptions to the Octet Rule The Incomplete Octet BeH2 BF3

  20. N – 5e- S – 6e- N O 6F – 42e- O – 6e- 48e- 11e- F 6 single bonds (6x2) = 12 F F Total = 48 S 18 lone pairs (18x2) = 36 F F F Exceptions to the Octet Rule Odd-Electron Molecules NO The Expanded Octet (central atom with principal quantum number n > 2) SF6

  21. Dative Covalent Bond / Coordinate Covalent Bond • A covalent bond in which one of the atoms donates both electrons. • Examples: - NH4+, NH3AlCl3, NH3BF3

  22. Hybridization Hybridization – mixing of two or more atomic orbitals to form a new set of hybrid orbitals. • Mix at least 2 nonequivalent atomic orbitals (e.g. s and p). Hybrid orbitals have very different shape from original atomic orbitals. • Number of hybrid orbitals is equal to number of pure atomic orbitals used in the hybridization process. • Covalent bonds are formed by: • Overlap of hybrid orbitals with atomic orbitals • Overlap of hybrid orbitals with other hybrid orbitals

  23. Formation of sp3 Hybrid Orbitals

  24. Formation of Covalent Bonds in CH4

  25. sp3-Hybridized N Atom in NH3

  26. Formation of sp Hybrid Orbitals

  27. Formation of sp2 Hybrid Orbitals

  28. How to predict the hybridization of the central atom? • Draw the Lewis structure of the molecule. • Count the number of lone pairs AND the number of atoms bonded to the central atom # of Lone Pairs + # of Bonded Atoms Hybridization Examples 2 sp BeCl2 3 sp2 BF3 4 sp3 CH4, NH3, H2O 5 sp3d PCl5 6 sp3d2 SF6

  29. Bonding in Ethylene, C2H4 Sigma bond (s) – covalent bonds formed by orbitals overlapping end –to-end , with the electron density between the nuclei of the bonding atoms Pi bond (p) – a covalent bond formed by sideways overlapping orbitals with electron density concentrated above and below plane of nuclei of the bonding atoms

  30. Another View of p Bonding in Ethylene, C2H4

  31. Bonding in Acetylene, C2H2

  32. H O C H Describe the bonding in CH2O C – 3 bonded atoms, 0 lone pairs C – sp2

  33. H H C H C O O H Sigma (s) and Pi Bonds (p) Single bond 1 sigma bond 1 sigma bond and 1 pi bond Double bond Triple bond 1 sigma bond and 2 pi bonds How many s and p bonds are in the acetic acid (vinegar) molecule CH3COOH? s bonds = 6 + 1 = 7 p bonds = 1

  34. Intermolecular Forces • Intermolecular forces: attractive forces between molecules. • Van der Waals forces: - the attractive or repulsive force between molecules due to covalent bonds or to the electrostatic interaction of ions with one another or with neutral molecules. • The term includes: - permanent dipole–permanent dipole forces - instantaneous induced dipole-induced dipole (London dispersion force). Examples: interaction between H2, Cl2, F2, CH4

  35. or … … H H B A A A Hydrogen Bond The hydrogen bond is a special dipole-dipole interaction between the hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom. A & B are N, O, or F

  36. Decreasing molar mass Decreasing boiling point Why is the hydrogen bond considered a “special” dipole-dipole interaction?

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