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Unit 8: Covalent Bonding

Unit 8: 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 8: Covalent Bonding

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  1. Unit 8: 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. Lewis Structures and Formal Charge For each atom in a molecule… • Formal charge = # valence e– – # lone e– – # bonds or = # valence e– – dots – lines • Sum of formal charges = overall charge on molecule/ion Formal Charges Water H2O H: 1 - 1 = 0 O: 6 - 4 - 2 = 0 Total charge = 0 Hydronium H3O+ H: 1 - 1 = 0 O: 6 - 2 - 3 = +1 Total charge = +1

  13. 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

  14. Selected Shapes and Geometries using VSEPR “Things”

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

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

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

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

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

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

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

  22. 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.

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

  24. 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:

  25. 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

  26. 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

  27. 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-

  28. 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

  29. 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

  30. - + + + 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.

  31. 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+

  32. 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

  33. 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–

  34. 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!

  35. 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.

  36. 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

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

  38. 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

  39. 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!

  40. 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

  41. 2p 2s E 1s Orbital Review • Each atom contains electrons in atomic orbitals Each orbital can hold a max of 2 electrons Valence e-  Core e- • Orbitals are in energy levels • Inner energy levels = core e- (kernel e-) • Outermost energy level = valence e- • Valence e- are used for bonding • Ionic bonding: transfer of those e- • Covalent bonding: sharing of those e-

  42. E 1s 2p E 2s 1s 90° 2p 1s 2p Orbitals and Bonding • Covalent bonds are formed by sharing two e- • Sharing occurs in overlapping atomic orbitals 1s 1s : • Example: H2 (H-H) • Use 1s orbitals for bonding • Example: H2O • From VSEPR: bent, 104.5° angle between H atoms • Use two 2p orbitals for bonding? How can we explain this bonding and match VSEPR?

  43. O=O or O2 2p 2p Bonding • Single bonds (one sigma bond) σ Overlap of s orbitals on bond axis • Double bonds (one sigma + one pi bond) Sharing of electrons between p orbitals perpendicular to the bonding atoms Termed “pi” or π bonds Bond Axis of σ bond • One π bond

  44. Hybrid Orbitals + = H2 2H: 1s + = LINEAR sp C: two sp C: 2s, 2p TRIGONAL PLANAR + = sp2 C: three sp2 C: 2s, 2px, 2py = + TETRAHEDRAL sp3 C: four sp3 C: 2s, 2px, 2py, 2pz

  45. O H H Hybrid Orbitals: Tetrahedral Mix starting atomic orbitals together to form hybrid orbitals • Bonds with H Lone pairs • 1 s and 3 p orbitals on O mix to form 4 identical sp3 hybrid orbitals (tetrahedron) • Two sp3 orbitals contain O lone pairs • Two form  bonds with H 1s orbitals

  46. Hybrid Orbitals: Trigonal Planar  bond  bonds • 1 s and 2 p orbitals on C mix to form 3 identical sp2 hybrid orbitals (trigonal planar). • One p orbital remains unhybridized. • sp2 orbitals form  bonds with O and H atoms • Leftover p orbital forms  bond with O

  47. Hybrid Orbitals: Linear Bonding  bonds  bonds • 1 s and 1 p orbital on C mix to form 2 identical sp hybrid orbitals (linear) • Two p orbitals remain unhybridized. • Each sp orbital forms a  bond with O • Each leftover p orbital forms a  bond with O

  48. Determining Hybridization • To determine an atom’s hybridization, count “things” around the atom “Things” = bonded atoms and lone pairs • Two things: sp hybrid orbitals • Three things: sp2 hybrid orbitals • Four things: sp3 hybrid orbitals Methane Acetonitrile sp sp3 4 “things” sp3 hybridization

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