Unit 5B: Covalent Bonding

1. Unit 5B: Covalent Bonding

2. Bonding Review Covalent Bonds (2 nonmetals) …atoms share e– to get a full valence shell C 1s2 2s2 2p2 F 1s2 2s2 2p5 *Both need 8 v.e – for a full outer shell (octet rule)!* 4 valence e- 7 valence e- o x x C x F o o x x x x o

3. Draw the Lewis dot structure for the following elements (write e- config first): Si O P B Ar Br 1s2 2s2 2p6 3s2 3p2 4 valence e- 1s2 2s2 2p4 6 valence e- 1s2 2s2 2p6 3s2 3p3 5 valence e- 3 valence e- 1s2 2s2 2p1 1s2 2s2 2p6 3s2 3p6 8 valence e- 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 7 valence e-

4. Notice any trends…? 1 2 3 4 5 6 7 8 H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Se Br Kr TRANSITION METALS Rb Sr Te I Xe Cs Ba The group # corresponds to the # of valence e–

5. F F C F F F F F F C F F F Let’s bond two F atoms together… Each F has 7 v. e– and each needs 1 more e– F2 F Now let’s bond C and F atoms together… carbon tetrafluoride (CF4)

6. Lewis Structures: 2D Structures NH3 CH2O CO2 SO2 CH4

7. Drawing Lewis Structures • Sum the # of valence electrons from all atoms Anions: add e– (CO32- : add 2 e– ) Cation: subtract e– (NH4+: minus 1 e– ) • Predict the arrangement of the atoms • Usually the first element is in the center (often C, never H) • Make a single bond (2 e–) between each pair of atoms • Arrange remaining e– to satisfy octets (8 e– around each) • Place electrons in pairs (lone pairs) • Too few? Form multiple bonds between atoms: double bond (4 e–) and triple bond (6 e–) • Check your structure! • All electrons have been used • All atoms have 8e- Exceptions: Remember that H only needs 2e– !

8. H C N Lewis Structure Practice Draw a Lewis Structure for the following compounds: • CH4 • H2O • NF3 • HBr • OF2 • HCN • NO3- • CO32-

9. Lewis Structure Trends Here are some useful trends… C group • Forms a combo of 4 bonds and no LP (Lone Pairs) • i.e. CO2 N group • Forms a combo of 3 bonds and 1 LP • i.e. NH3 O group • Forms a combo of 2 bonds and 2 LP • i.e. CH2O F group (halogens) • Forms 1 bond and 3 LP • i.e. OF2 Note that these are NOT always true!

10. Carbonite Carbonate? CO32- CO22-

11. Resonance Structures Show resonance Show movement of e- Resonance structures differ only in the position of the electrons • The actual structure is a hybrid (average) of the resonance structures • Technically NOT two single bonds and one double bond • All 3 Oxygen atoms share the double bond • 3 equal bonds (somewhere between a double and single) • Arrow formalism: curved arrows show electron movement

12. Predicting Molecular Shape: VSEPR (Valence Shell Electron Pair Repulsion) • Electrons repel each other • The molecule adopts a 3-D shape to keep the electrons (lone pairs and bonded e-) as far apart as possible • Different arrangements of bonds/lone pairs result in different shapes • Shapes depend on # of bonds/lone pairs (“things”) and LP around the central atom

13. Selected Shapes and Geometries using VSEPR “Things”

14. Carbon Dioxide: CO2 O C O Lewis Structure • Two “things” (bonds or lone pairs) • Linear geometry • 0 LP → Linear Shape • 180o Bond angle

15. O C H H Formaldehyde: CH2O Lewis Structure • Three “things” • Trigonal planar geometry • 0 LP → Trigonal planar shape • 120° bond angles

16. A S B O O A A Sulfur Dioxide: SO2 Lewis Structure • Three “things” • Trigonal planar geometry • 1 LP → Bent shape • 120° bond angles

17. Methane: CH4 Lewis Structure • Four “things” (bonds/LP) • Tetrahedral geometry • 0 LP → Tetrahedral shape • 109.5o bond angles

18. Ammonia: NH3 Lewis Structure • Four “things” (bonds/LP) • Tetrahedral geometry • 1 LP → Trigonal pyramid shape • 107o bond angles

19. Water: H2O Lewis Structure • 4 “things” (bonds/LP) • Tetrahedral Geometry • 2 LP → Bent Shape • 104.5o bond angle

20. Hydrogen Chloride: HCl H Cl Cl Lewis Structure • Four “things” (bonds/LP) • Tetrahedral geometry • 3 LP → Linear Shape • NoBond angle

21. A special note… H Cl N O Br Cl N O For any molecule having only two atoms… • e.g. N2, CO, O2, Cl2, HBr, etc. • Geometry = Linear • Shape = Linear • Bond Angle(s)? = None • It is much like geometry… what is formed by connecting two points? …a line.

22. You will need to commit these to memory! “Things”

23. VSEPR Practice (w/o aid of yellow sheet) • CO2 G: S: Angle: • ClO2- G: S: Angle: • NO2- G: S: Angle: • CH3COO- G: S: Angle: • PBr3 G: S: Angle: • AsO43- G: S: Angle:

24. Electronegativity and Bond Type The electronegativity difference between two elements helps predict what kind of bond they will form. Definition e- are evenly shared e- are unevenly shared e- are exchanged (gained or lost) Electronegativity difference ≤ 0.4  0.5 – 1.8  > 1.8 Bond type Covalent  Polar covalent  Ionic

25. H 2.1 Li 1.0 Be 1.5 B 2.0 C 2.5 N 3.0 O 3.5 F 4.0 Na 0.9 Mg 1.2 Al 1.5 Si 1.8 P 2.1 S 2.5 Cl 3.0 K 0.8 Ca 1.0 Br 2.8 I 2.5 Practice with Bond Types Bond Type? Ionic Covalent Polar covalent Covalent Sample Bonds NaCl Cl-Cl C-O C-H Electronegativity Difference 3.0 – 0.9 = 2.1 3.0 – 3.0 = 0 3.5 – 2.5 = 1.0 2.5 – 2.1 = 0.4 Electronegativity difference ≤ 0.4 0.5 – 1.8  > 1.8 Bond type Covalent  Polar covalent  Ionic

26. Dipole Moments and Polarity • Occurs in polar covalent bonds • Uneven distribution of e- • Atoms become partially charged Partially “+” charged end Arrow points toward partially “-” end δ+ δ-

27. HCN CO2 CO32- CH2O SO2 CH4 CH3F C3H8 CO NH3 Polarity Examples • Check molecule for dipole moments (polar bonds) • When determining overall polarity, an imbalanced structure will likely be polar (at least partially) • Even with polar bonds, a balanced structure is non-polar overall • Any structure with lone pairs on the central atom is automatically polar! Try these with your neighbors… Non-polar Polar Polar Non-polar Non-polar Non-polar Polar Polar Polar Polar