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Unit 5B: Covalent Bonding

Unit 5B: Covalent Bonding. Bonding Review. Covalent Bonds (2 nonmetals) …atoms share e– to get a full valence shell C 1s 2 2s 2 2p 2 F 1s 2 2s 2 2p 5 *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.

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Unit 5B: Covalent Bonding

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  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) • e.g. CO2 N group • Forms a combo of 3 bonds and 1 LP • e.g. NH3 O group • Forms a combo of 2 bonds and 2 LP • e.g. CH2O F group (halogens) • Forms 1 bond and 3 LP • e.g. 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

  28. Intermolecular Forces (IMF’s) • Intramolecular Forces = bonding within a molecule • e.g. ionic, covalent, polar covalent bonds • Intermolecular Forces = interactions between two molecules • …Intercity v. Intracity v. Innercity • Intermolecular Forcesare ALL weaker than Intramolecular bonds

  29. - + + + Na+ Na+ Na+ Cl– Cl– Cl– IMF’s: Ion-Ion Force Opposite Charges Attract Similar Charges Repel Attractive and repulsive forces between two separate ions.

  30. H H H Cl H Na+ Na+ IMF’s: Ion-Dipole Force The interaction between an ion and another molecule that has a dipole moment. (polar covalent) δ+ Cl + + δ- δ- δ+ Lewis Structure δ- O δ+ δ+

  31. H H H H Cl Cl IMF’s: Dipole-Dipole Force δ+ δ- The interaction between two separate molecules, each having a dipole moment. (polar covalent) Cl Cl δ- δ+ HCl = Stomach Acid

  32. H H H H O O IMF’s: Hydrogen Bonding A specific type of dipole-dipole interaction between an H bond donor and an H bond acceptor. H bond donor: an H bonded to N, O, or F H bond acceptor: any lone pair of e–

  33. H H H H H H H H H H H H H H H H IMF’s: London Dispersion Forces Involves an instantaneous dipole. This dipole will induce dipoles in other molecules. Probable? NO! Possible? YES! Probable? Yes Possible? Yes Probable? Yes Possible? Yes Probable? Yes Possible? Yes δ+ δ- Why instantaneous? This dipole will only remain for an instant! The electrons will quickly move to another part of the molecule! δ- δ+ All molecules will exhibit LDF ↑ mass,↑ LDF Instantaneous = WEAKEST!

  34. IMF Review STRONGEST Ion-Ion Ion-Dipole Hydrogen Bonding Dipole-Dipole London Dispersion Forces (LDFs) • a.k.a. van der Waals Forces Involves an ion (+ and – charged) Involves a dipole (polar molecule) Weakest Involves a non-polar molecule *Remember: These are all weaker than actual bonds (ionic, covalent, etc.). These are just attractions.

  35. O C H H IMF Practice Formaldehyde: CH2O Lewis Structure Trigonal planar geometry 120° bond angles Polar C=O bond = Net dipole moment IMF= Dipole-Dipole

  36. Methane: CH4 Lewis Structure Tetrahedral geometry 109o bond angles Covalent bonds = No net dipole IMF = London dispersion forces

  37. Ammonia: NH3 Lewis Structure Trigonal pyramid 107o bond angles Polar Bonds, Lone pairs = Dipole 1 H bond acceptor (LP), 3 H bond donors (N-H) IMF = Hydrogen Bonding

  38. O C O Carbon Dioxide: CO2 Lewis Structure Linear geometry 180o Bond angle C=O bond is polar, but… Dipoles cancel = No net dipole IMF = London Dispersion Forces!

  39. Water: H2O Lewis Structure Bent Polar bonds, lone pairs = Net dipole 2 H bond acceptor (LP), 2 H bond donors (O-H) IMF = Hydrogen bonding

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