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Ch. 16 Covalent Bonding

Ch. 16 Covalent Bonding. VSEPR Theory, Polarity, and using Electronegativity. Covalent Bonds. Forms when 2 atoms share a pair of valence e - A. Types of Covalent Bonds 1. Single Covalent Bond – two atoms share one pair of electrons Ex: F 2.

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Ch. 16 Covalent Bonding

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  1. Ch. 16 Covalent Bonding VSEPR Theory, Polarity, and using Electronegativity

  2. Covalent Bonds • Forms when 2 atoms share a pair of valence e- A. Types of Covalent Bonds 1. Single Covalent Bond – two atoms share one pair of electrons Ex: F2 Unshared pair – e- not shared between atoms ● ● ● ● ● ● ● ● ● ● ● ● F F F F ● ● ● ● ● ● F F ● ● ● ● ● ● ● ● or ● ● ● ● ● ● ● ● ● ● ● ● ● ● What makes this bonding work? Atoms have 8 e- in their outer level to make them stable

  3. Covalent Bonds (cont.) Ex: H2 H H H ● H H H or ● ● ● Why does H2 only need 2 e- to be stable? first energy level only contains 2 e-

  4. ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● Covalent Bonds (cont.) 2. Double Covalent Bond – 2 pairs of electrons are shared between atoms Ex: O2 ● ● ● ● O O O O O O ● ● ● ● ● ● or ● ● ● ● ● ●

  5. Covalent Bonds (cont.) 3. Triple Covalent Bond – 3 pairs of electrons are shared between atoms Ex: N2 ● ● ● ● N N N ● N N N ● ● ● ● ● ● ● ● ● ● ● or ● ● ● ● ● ● ● ●

  6. Covalent Lewis Dot Structures 1. Determine the # of valence e- in each atom in the molecule (# valence e- = roman numeral for group A atoms) 2. The central atom is often the first atom written & is usually the atom with the least # of e-. (Exception – H can’t be the central atom). This is going to be the atom that needs to share the most electrons.

  7. ● ● Lewis Dot Structures for Compounds 3. Place the electrons around the atoms so each is stable (8 around it, except H – only 2) Examples: 1. Br2 ● ● ● ● ● ● Br Br Br Br ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ●

  8. ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● 2. NH3 ● H N N H ● ● ● H H ● H ● H 3. CO2 O ● ● C ● ● ● O C O ● O ● ●

  9. ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● Cl 4. CCl4 ● Cl Cl ● C ● C Cl Cl ● ● Cl ● ● Cl Cl ● 5. H2O ● H H O O ● ● H H ●

  10. Covalent Bond Practice Problems: 1. CH4 4. OF2 2. H2 5. CHI3 3. PH3 6. CO2

  11. VSEPR Theory • Explains the shapes of molecules. • The VSEPR theory states: b/c electrons repel each other, molecules adjust their shapes so that the valence e- pairs are as far apart from each other as possible.

  12. Bond Polarity Polar Covalent Bond – when 2 atoms are joined by a covalent bond and the bonding electrons are not shared equally

  13. Bond Polarity (cont.) Nonpolar Covalent Bond – when 2 atoms are joined by a covalent bond and the bonding electrons are shared equally

  14. Differences between polar, nonpolar, and ionic bonds

  15. How do you determine if a bond is polar, nonpolar, or ionic? Subtract the electronegativities of the bonding atoms (p. 405 in textbook)

  16. Electronegativity Differences & Bond Type

  17. Tell if the bonds between the following atoms are polar, nonpolar, or ionic: 1. Hydrogen and Carbon 2. Oxygen and Carbon 3. Potassium and Chlorine 4. Fluorine and Fluorine 5. Nitrogen and Oxygen H 2.1 C 2.5 0.4 Nonpolar O 3.5 C 2.5 1.0 Polar K 0.8 Cl 3.0 2.2 Ionic F 4.0 F 4.0 0.0 Nonpolar N 3.0 O 3.5 0.5 Polar

  18. Polarity of Molecule Polar Molecule – a molecule with a positive and negative end. Polar bonds must be present.

  19. Polarity of Molecule (cont.) • It is possible to have polar bonds but not a polar molecule! • Carbon dioxide has 2 polar bonds and is linear. • Bond polarities cancel out b/c they are in opposite directions. Carbon Oxygen Oxygen

  20. Practice: Write the dot structure of the following molecules – then predict the shape and polarity • I2 • PCl3 • H2S • CHI3 • SiO2 • CH2O

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