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

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Presentation on theme: "Ch. 16 Covalent Bonding VSEPR Theory, Polarity, and using Electronegativity."— Presentation transcript:

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: F 2 FF ●● ● ● ● ● ●● ● ● ● ●●● FF ●● ● ● ● ● ●● ● ●● ●●● or FF ●● ● ● ● ●● ● ● ●●● Unshared pair – e- not shared between atoms What makes this bonding work? Atoms have 8 e- in their outer level to make them stable

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

4 2. Double Covalent Bond – 2 pairs of electrons are shared between atoms Ex: O 2 O ●● ●● ● ● ● O ●● ●● ● ● O ●● ● O or ● ● ● ● ●● ●● O ●● O ●● ●● ● ●

5 Covalent Bonds (cont.) 3. Triple Covalent Bond – 3 pairs of electrons are shared between atoms Ex: N 2 NN ●● ●● ● N ●● ●● ● N ● ● N ● ● or N ● ● ● ● ● ● ● ● ●●

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. Br 2 Br ●● ● ● ● ● ● ●● ● ● ● ● ● ●● ● ● ● ● ●● ● ● ● ●

8 2. NH 3 N ● ●● ● ● H ● H ● H ● N ●● H H H 3. CO 2 C ● ● ● ● O O ●● ●● ●●●● ● ● ● ● COO ●● ●● ●● ●●

9 4. CCl 4 C ● ● ● Cl ● ● ● C 5. H 2 O H ● O ●● ●● ● ● ● Cl ●● ●● ●● ●● ●● ●● ●● ●●●●●●●● ● ●● ●● ●●●● ●● ●● ●● ●● ●● ●●●●●●●● H ● O H H ●● ●●

10 Covalent Bond Practice Problems: 1. CH 4 4. OF 2 2. H 2 5. CHI 3 3. PH 3 6. CO 2

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 ShapeFormulaBond AngleElectrons Linear AX 2 180 o 4 shared 0 unshared Linear AX180 o 1 shared 3 unshared Bent AX 2 105 o 2 shared 2 unshared Trigonal Pyramidal AX 3 107 o 3 shared 1 unshared Tetrahedral AX 4 109.5 o 4 shared 0 unshared Trigonal Planar AX 3 120 o 4 shared 0 unshared Contains a double bond

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

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

15 Differences between polar, nonpolar, and ionic bonds

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

17 Electronegativity Differences & Bond Type Type of Bond Electronegativity Difference Range Nonpolar Covalent Bond 0.0 – 0.4 Polar Covalent Bond 0.5 – 2.0 Ionic Bond greater than 2.0

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

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

20 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

21 Practice: Write the dot structure of the following molecules – then predict the shape and polarity 1.I 2 2.PCl 3 3.H 2 S 4.CHI 3 5.SiO 2 6.CH 2 O


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