3. Triple covalent bonds include three shared pairs of electrons. Example N2.

15.3 Covalent Compounds

1. Covalent molecules can be written as electron dot formulas (Lewis Structures): try water ammonia, methane, and carbon dioxide. (See page 397-398)

2. Water:

3. Ammonia:

4. Methane:

4. Carbon usually forms 4 covalent bonds instead of 2. One 2s electron is moved to an empty 2p orbital.

6. Small amount of energy to bring electron up. CH₄ is much more stable than CH₂.

5. Do practice problem 8. a., b., c.

6. Carbon dioxide:

7. Figure 15.7, CO2

15.4 Coordinate Covalent Bonds

1. A coordinate covalent is formed when one atom contributes both bonding electrons in a covalent bond. Example: Carbon monoxide (see page 400).

2. If double bond (C=0), Oxygen is stable, but carbon is not. If oxygen donates a PAIR of unshared electrons then both are stable C=O.

3. Ammonium, NH₄+, has a coordinate covalent bond. There are 3 covalent bonds from the original Ammonia atom. H+ is attracted to and shares BOTH unshared electrons of N

4. See Example 2, page 401 for H₃0+

5. Example 3,polyatomic anion SO₃ ^2-

6. Do practice problem 10.

15.5 Resonance:

1. Resonance structures occur when two or more valid electron dot formulas can be written for a molecule.

2. Example:

Double headed arrows used to show resonance.

3. Molecules are really a hybrid of the forms. Characteristics are in between.

4. Resonance strucutres have some characteristics of its resonance forms, but it is a disctinct species.

15.6 Exceptions to the octet rule - Honors (Octet rule applies mainly to 2nd-period elements).

1. There are three categories of exceptions:

2. The incomplete octet:

E.g., Be in period 2: BeH₂ (H-Be-H). Be can only have 4 valence electrons.

E.g., B and Al in period 3 - BF₃ is deficient by 2 electrons. It readily reacts with ammonia to make BF₃^.N₃. Forms coordinate covalent bond.

3. Odd-Electron Molecules:

If odd number of electrons in molecules, then impossible to write dot structures for octet rule.

4. Example NO₂. (also shows resonance)

5. An unpaired electron is always present. One oxygen has only 7 valence electrons.

6. Expanded Octet: In some compounds there are more than 8 electrons around central atom. In and beyond 3rd period. In addtion to 3s, 3p, 3d are involved in bonding.

7. P and S sometimes go to 10 or 12 electrons. Phophorus trichloride (PCl₃) and Phosphorus pentachloride (PCl₅).

PCl₃ follows octet rule. PCl₅ dosen't. P has 10 valence electrons.

8. SF₆ - S has 12 valence electrons

9.Electrons are spinning electric charges and create magnetic fields.

5. Diamagnetc is when all electrons are paired. Opposite spins cancel out. Only weakly repelled by external magnetic field.

5. Paramagnetic is when substances contain one or more unpaired electrons. They show a relatively strong attraction to an external magnetic field.

In the presence of a magnetic field paramagnetic substances are heavier.

6. Ferromagnetic is a stronger attraction by Fe²⁺, Co²⁺, Ni²⁺. Ions with unpaired electrons line up with magnetic field and stay that way. Large groups of ions in the metals.

7. Oxygen (O₂) is paramagnetic.

8. This is possible if there are unpaired electrons. Experimental evidence (distance of bonds) indicates resonance:

15.7 Molecular Models - (Molecular Orbital Theory) Honors

1. Just as there is quantum theory of atomic orbitals, there is quantum theory of molecular orbitals.

2. When two atoms combine, their atomic orbitals overlap to produce molecular orbitals.

3. A molecular orbital belongs to the molecule as a whole.

4. 2 electrons are required to form a molecular orbital.

5. Overlap of 2 half filled atomic orbitals produces filled single covalent bond.

6. Molecular Orbital Model requirels that number of molecular orbitals equal number of atomic orbitals.

7. Two overlapping atomic orbitals produce two molecular orbitals, though only one bond (single covalent).

Bonding orbital - a molecular orbital whose energy is lower than that of the atomic orbitals from which it is formed.

Antibonding orbital - a molecular orbital whose energy is higher than that of the atomic orbitals from which it is formed.

8. Illustration

9. Hydrogen molecule (H₂) has 2 overlapping atomic orbitals. 1s electrons are shared.

10. Illustration

11. Electrons seek the lowest energy.

12. Both electrons fill the bonding molecular energy level, making stable covalent bond.

13. In this level the probability is finding electrons between nuclii.

14. A sigma bond (s ) is formed when two atomic orbitals combine to form a molecular orbital that is symmetrical along the axis connecting two atomic nuclei.

15. Illustration

16. There is always a balance between attraction of opposites and repulsion of like charges.

17. There is an attraction of hydrogen nuclei to the electrons between them in bonding orbital.

18. Other cases bonding results in electron pairs in higher level. Electrons are not between nuclei.

19. In a theoretical He₂ molecule, 2 electrons would go into the bonding orbital, and 2 would go into the antibonding orbital. Result is repulsions make it unstable. He remains an atom.

20. P orbitals also combine. E.g., F₂ has 2 half filled p orbitals which overlap end to end, producing a s bond along the axis between the 2 nuclei.

21. Illustration

22. P orbitals may also overlap side by side, producing a pi bond (P ). In a Pbond the bonding electrons are most likely to be found in a sausage shaped regions above and below the bond axis of the bonded atoms.

23. Illustration

24. Orbital overlap in P is not as extensive and is weaker than s .

15.8 VSEPR THEORY (Valence-shell electron-pair repulsion model) Chang - Molecular Geometry.

1. Properties, such as mp, bp,density, reactions, etc. are affected by geometry.

2. Electron dot strucures don't show 3 dimensions. Example: methane CH₃ is a tetrahedral H-C-H angle of 109.5^o

3. VSEPR stands for valence-shell electron-pair repulsion theory.

4. VSEPR theory states that because the electron pairs repel, molecules adjust their shapes so that the valence-electron pairs are as far apart as possible.

6. Can predict geometry in a systematic way. Divide molecules into two categories: central atom has lone pairs or not.

7. If no non-bonding pairs then forms in Table 10.1 (Chang)

8. Molecules have bonding and lone pairs. Order of repulsion

lone vs lone > Lone vs bonding > bonding vs bonding.

9. Lone pairs have greater spatial distribution and greater repulsion.

10. Number of lone pairs affects geometry of molecule.

11. Figure 10.2 Chang
 

Figure 10.3 (combination)

12. Ammonia (NH₃) is pyramidal. Unshared pair of electrons is held closer than the bonding pairs and repels bonding pairs to angle of 107^o (experimentally determined)

13. In water there are 2 bonding pairs and 2 unshared pairs. This forms a tetrahedral but planar molecule. Angle = 105^o. H's repelled by unshared pairs.

14. CO₂ has no unshared pairs so is linear. O=C=C. 180^o.

15. Figure 15.19

16. Know example angles.

15.9 Hybrid Orbitals - Honors

1. Hybridization is a description of molecules that takes both bonding and shape into account.

2. Hybridization - several atomic orbitals mix to form the same number of equivalent hybrid orbitals.

3. Single covalent bond - methane: one 2s orbital and three 2p orbitals of C mix to form four sp3 hybrid orbitals. Angle = 109.5^o .

4. Four sp3 orbitals of C overlap with the 1s orbitals of H atoms. Since sp3 extend farthe and overlap more, the s bonds are very strong. 8 e- in bonds fill octet of C and duet of H.

5. Molecule is tetrahedral. (no lone pairs)

5. Double covalent bond - Ethene: H₂C=CH_2, one 2s orbital and 2 2p2 orbitals combine to form 3 2sp2 orbitals, with angle of 120^o between HCH. One 2p orbital is left.

6. On each C 2 sp3 orbitals form s bonds with H's. 3rd sp3 orbitals form s bond between C's.

7. A second bond is formed between C with remaining nonhybridized 2p orbital, forming weaker P bond.

6. Figure 15.21

7. Ethene held together with 5 s and 1 P bond. Octet and duets are complete.

(both sigma and pi bonds have pair of electrons). Molecule is planar.

8. Tripple covalent bond - Ethyne (acetylene): 2s orbital combines with a 2p orbital to form 2 2sp orbitals. 2 p2 non hybridized orbitals remain.

9. One sp orbital combines C-C.

10. The other combines with a H to make a linear molecule.

11. Figure 15.22

15.10 POLAR BONDS

1. The nature of a covalent bond depends on the number and kinds of atoms joined. Not all bonds are the same. They will determine properties of molecule.

2. Nonpolar covalent bond is when atoms are the same and bonding electrons are shared equally. Examples, O₂, H₂, N₂

3. Polar covalent bond is when two different atoms are joined by a covalent bond and the bonding electrons are shared unequally.

4. The higher the electronegativity (page 366), the stronger attraction for electron pairs.

5. This gives some atoms a relative stronger charge. Example H-Cl. Cl is more electronegative (3.0 vs 2.1). So molecule of HCl is slightly negative on the Cl side (shared electrons are closer to Cl). H is slighty more positive. Use greek delta. (d ) 417.

6. Example - water (page 418) is polar.

7. Electronegativity will determine type of bond. 418. Difference between electronegativities will determine bond. If greater than 2.0 then ionic.

8. Table 15.3

15.11 POLAR MOLECULES

1. Polar molecule is when one end of the molecule is slightly negative, and one end is slightly positive.

2. A molecule with with 2 poles is called dipolar or dipol. Example HCl.

3. Not every molecule with polar bonds is itself polar. Example CO₂ or CCl₄ where molecules are symmetrical and polarities cancel.

4. Do #27.

5. The polarity of the entire molecule depends on the shape of the molecule.

6. Since water is bent it is polar.

15.12 Bond Dissociation Energies

1. A large quantity of heat is liberated when hydrogen atoms combine to form hydrogen molecules.

2. Product is more stable than reactants.

3. 435kJ of energy is required to dissociate 1 mol of hydrogen.

4. The energy required to break a single bond is known as bond dissociation-energy.

5. See table 15.4 for some dissociation energies.

6. Higher the energy required, the more stable and less likely to react.

15.13 Intermolecular Attractions

1. In addition to attraction within molecules, there are attractions between molecules.

2. Though weaker than covalent bonds, they determine whether a compound is a solid, liquid, or gas.

3. The weakest attractions between molecules are collectively called van der Waals forces ( Johannes van der Waals, 1837-1923). There are 2 major forces:

Dispersion forces, the weakest of all molecular interactions, are thought to be caused by the motion of electrons. Strength of forces increase with increase in # of electrons

Halogens show this. F and Cl with few electrons are gases at STP. Br is liquid and I is a solid at STP.

Dipole interactions, occur when polar molecules are attracted to one another. Electrostatic attractions occur between oppositely charged regions of dipolar molecules.

E.g., since water is a polar molecule, the H side of one molecule is attracted to the O side of another. Figure 15.28 In the case of water the relatively strong interaction is called a hydrogen bond.

Hydrogen bonds are attractive forces in which hydgrogen that is covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of an electronegative atom in the same molecule or in a nearby molecule.

Hydrogen is only element whose electrons are not shielded from nucleus. A highly polar bond forms when it bonds to an electronegative atom like O, N, or F. Since H electron is drawn to electronegative atom, shares and unbonded pair from another molecule.

Though only 5% attraction of covalent bond, it is stronges intermoleclar force. Examples from Prob. 30

15.14 Properties of Molecular Substances

1. Properties of compounds depends on types of bonds: solids, liquids, gases.

2. See table 15.5.

3. Intermolecular attractions effect properties of covalent compounds.

4. Most are low mp and bp. Few are very high. Mp = 1000 - 3000oC.

5. Network solids (or crystals) are stable substances in which all of the atoms are covalently bonded to each other.

6. E.g., diamonds (no melting point, but vaporizes at 3500) and silicon carbide (mp = 2700oC).

7. Substances are thought of as large single molecules.

Suplement:

I. Formal Charge: #electrons in atom - nobonded pairs -¹/₂bonded pairs = charge of atom/ion

II. Valence Bond Theory (Chang, 387)

1. 1st of 2 quantum mechanical theories - states that electrons in molecule occupy atomic orbitals of the individual atoms.

2. In a H-H bond there is an overlap in space between the 2 atoms.

3. When far apart in space, 2 atoms have 0 PE and no intereaction.

4. As atoms approach e- attract nuclei of other atoms. As get closer, e- and nuc.+ also repel each other. While separated, attraction is stronger than repulsion.

5. As get closer, PE decreases till it reaches minimum. This is most stable and is overlap of 1s and is H₂.

6. Law of C of E: As PE decreases to form H₂, heat must be given off. Converse. To break H₂ requires E.

7. VB theory accounts for difference between bonds in energy and lengths, while Lewis doesn't.

III. Dipole Moments

1. There is a shift in electron density from one atom to atom of greater electronegativity.

2. This creates polar molecules.

3. A quantitative measure of the polarity is the dipole moment (m ).

Q = charge, r = distance between charges. m is always +.

4. Dipole moment is expressed in Debye units (D).

5. Diatomic molecules of same atom are nonpolar and have no dipol moments.

6. Diatomic molecules of different elements (HCl, CO, NO) do.

7. 3 or more elements depend on polarity of molecule AND geometry.

8. Use arrows to indicate BOND MOMENTS +®

9. The bond moment is a vector quantity and has magnitude and direction.

10. Example 1: NH3 and NF3

11. Example 2: C2H2Cl2

IV. Molecular Orbital Configurations

1. Must arrange molecular orbitals in order of increasing energy.

2. # molecular orbitals = # of atomic orbitals.

3. Bond Order = 1/2(#e- in bonding Mos - # e- in antibonding MOs)

4. Molecular orbital energy level diagram: must place total molecular electrons in bonding and antibonding orbitals.

5. Molecular orbital electron configuration: fill sigma and pi orbitals in order of relative energy.

6. Relative Energy of Molecular Bonds