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Lecture 9: Covalent Bonding: Part 1 Lewis Structures (Ch. 7)

Lecture 9: Covalent Bonding: Part 1 Lewis Structures (Ch. 7). Dr. Harris Suggested HW: 2, 7, 9, 13, 14, 17, 27d. . Recap. Ionic compounds are formed by the transfer of electrons from a metal to a nonmetal

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Lecture 9: Covalent Bonding: Part 1 Lewis Structures (Ch. 7)

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  1. Lecture 9: Covalent Bonding: Part 1 Lewis Structures (Ch. 7) Dr. Harris Suggested HW: 2, 7, 9, 13, 14, 17, 27d.

  2. Recap • Ionic compounds are formed by the transfer of electrons from a metal to a nonmetal • The result is a cation-anion pair. Each ion has reached its nearest noble gas configuration Na ([Ne] 3s1) + Cl ([Ne] 3s2 3p5) ---> Na+ Cl- Cl- Na+ [Ar] [Ne] • The Coulombic attraction between the oppositely charged ions is what holds the molecule together. This is why ionic compounds are solids at room temperature.

  3. Covalent Bonding • Covalent bonding is electron sharing (not transfer) between nonmetals • Consider Cl2(g) • Each Cl has 7 valence electrons • Chlorine is 1 electron short of a full octet ([Ar] configuration) • Since both atoms are the same, they have the same electron affinities and ionization energies, so 1 chlorine will not donate an electron to the other. [Ne]3s23p5 [Ne]3s23p5 Cl Cl

  4. Covalent Bonding • For both atoms to achieve an [Ar] configuration (full octet), they share a pair of electrons between them. An element in a covalent bond will react so that eight electrons occupy its valence shell. This is the OCTET RULE. • To indicate the covalent bond, we use a solid line. This is a single bond (2 electrons). Bonding represents an overlap of the valence orbitals. • The electrons not involved in bonding are called lone pairs. [Ar] [Ar] Covalent bond Cl Cl Cl Cl

  5. Covalent Bonding • All elements in a covalent bond will achieve an octet • Ex. OF2 • Oxygen has 6 valence electrons (needs 2) • Fluorine has 7 (each F needs 1) • To achieve octets, each F will share its electron with O [He]2s22p5 [He]2s22p5 [He]2s22p4 O F F F O F F O F [Ne] [Ne] [Ne]

  6. Electronegativity • Before we learn about how to draw a Lewis structure, it is very important to consider what makes ionic and covalent bonds so different: electronegativity • Electronegativity is the ability of an atom to attract electrons to itself. • Therefore, electrons are drawn more towards the more electronegative atom in a molecule. • Thus, when there is a difference in electronegativities between atoms in a molecule, the electrons are NOT equally shared.

  7. Table of Electronegativities Electronegativity of atoms increases up and to the right.

  8. Electronegativity • Although we have discussed ionic and covalent bonds, most chemical bonds are neither purely ionic or purely covalent • Most compounds are an intermediate between the two.

  9. Electronegativity • Let’s consider NaCl. The difference in electronegativity between Na and Cl is: • Because the difference in electronegativity is so big, the Na electron is completely pulled away by the Cl atom. So, the molecule is totally polar (ionic) • This is why there is no actual bond in ionic compounds, only coulombic attraction. 3.16 – 0.93 = 2.23 ionic Cl- Na+

  10. Electronegativity • Let’s look at another molecule, like HCl: • The HClmolecule is not purely ionic or covalent, but BOTH. • The electron density is unevenly distributed, such that more of the electron density is on the Cl than on the H. • Therefore, the H and Cl have partial charges. The arrow depicts the direction of electron “pull”, or dipole. Any molecule with a net dipole has polarity. We will discuss dipoles later. 3.16 – 2.1 = 1.06 polar covalent - + δ δ Cl H Partial positive character Partial negative character

  11. General Rules of Drawing Covalent Lewis Structures • Arrange atoms together. If possible, put the least electronegative atom in the center. Hydrogen is an exception to this rule. Hydrogen atoms are always terminal. • Compute the total number of valence electrons. Account for charges. • Represent bonds with solid lines. Ensure octets around all atoms, except hydrogen. Hydrogens can only accommodate 2 electrons. • Add in remaining lone pairs • If there are not enough electrons to complete an octet, consider multiple bonds or formal charges.

  12. Example • Draw the Lewis structure of CCl4 • Carbon is the least electronegative atom (also, there is only one C), so C is the central atom • Compute the valence electrons • C: [He] 2s2 2p2 Cl: [Ne] 3s2 3p5 C • Each Cl needs 1, the C needs 4. Therefore, the C will share each of its 4 valence electrons with Cl C Cl Cl Cl Cl Cl Cl Cl Cl Cl Cl C Cl Cl

  13. Hydrogen • Hydrogen CAN ONLY ACCOMMODATE 2 ELECTRONS • Ex. H2(g) [He]-like core • HX where X = F, Cl, Br, I (halogens) • In chemical structures, hydrogens are always terminal atoms, meaning that they are at the ends of a molecule H H H X

  14. Organic Molecules • Ex. Methanol, CH3OH Hydrogens are terminal, which means there is a C—O bond in the center C C C H H H H H O O O Now, C needs 3 more electrons around it to achieve an octet. O needs 1. H H H H C O H H H

  15. Double and Triple Bonds • There are instances where single bonds (2 e- bonds) are not enough to satisfy the octet rule. • Ex. C2H4 (ethene) C C C C C C • As you can see, C only has 7 electrons around it, including a lone electron on each C. This is highly unfavorable. • Lone electrons migrate into the C-C bond to form a 4e-double bond. H H H H H H H H H H H H H H H H H H C C C C H H

  16. Double and Triple Bonds • 6e-triple bonds are also possible, as in the case of N2 (g) N N N N N N Two unpaired electrons each N N 6 electron Triple Bond • The N atoms are sharing 6 electrons, so each N now has a full octet.

  17. Group Examples • Draw the following Lewis Structures • CH3Cl • NH3 • H2O • CO2 • HCN

  18. Formal Charges • So far, all of the atoms we’ve seen in covalent compounds have had a zero charge. • However, nonmetals can assume positive or negative charges in order to facilitate the formation of a covalent bond • Thus, to satisfy the octet rule in certain cases, formal chargeshave to be applied

  19. Formal Charges - As is, N can only make 3 bonds. X N N N H H H H H H H H H H H H - N loses an electron to accommodate the 4th H, attains a charge of 1+. H + + + N N H H H Consider ammonium, NH4+ -Lets start by looking at the valence electrons of the neutral atoms

  20. Formal Charge Allows Us to Rule Out Structures • Formal charges can help us determine if we’ve arranged a molecule incorrectly The preferred arrangement of atoms in a molecule is always the arrangement that requires the least amount of formal charge • If formal charges must exist, the more electronegative elements prefer negative formal charges, and less electronegative elements prefer positive ones.

  21. Ambiguity in Atomic Arrangement • Ex. Hydrogen cyanide, HCN • We know that C is the central atom in HCN. But how can we rule out HNC? Lets consider both possibilities 8 total 2 total 8 total • As we determined earlier, we have a CN triple bond and an HC bond. All octets are satisfied. • There are no formal charges required H C N

  22. Ambiguity in Atomic Arrangement • If we rearrange N and C, we have the following structure. As shown, C is 1 electron short of an octet. • N can not make anymore bonds because it already has an octet. N has a 1+ formal charge because it loses a valence electron. 7 total 2 total 8 total + H H N N C C 8 total 8 valence 2 valence • C must take on a 1- formal charge to reach an octet by accepting the N electron. Now, all the atoms are satisfied, and the molecule is neutral. - +

  23. Ambiguity in Atomic Arrangement • Now, our two possibilities are: • Both structures satisfy the octet rule, and both are charge neutral, but the HNC structure has formal charges on both N and C. Thus, the HCN structure is preferred. 8 total 2 total 8 total H H N C C N 8 total 8 total 2 total - +

  24. Group Example • The correct structure of the NOCl is a central N with a double bond to O, a single bond to Cl, and a lone pair on the N. • Use formal charges to show why O is notthe central atom.

  25. Resonance • When a covalent molecule has an overall charge, the charge is said to be delocalized, meaning that the charge is spread over the whole molecule, as opposed to being localized on a single atom • Ex. Nitrite NO2- • To draw this molecule, we put N as the central atom. We satisfy an octet on one O by making a N=O double bond. We can apply the 1- charge to the other O to complete its octet (O is more electronegative than N, so it will not go on N) O N O - The blue electron represents the electron added by applying the formal charge

  26. Resonance • You can’t distinguish between one O and the other. So, you could write either of the following: O N O O N O - - - O N O • These structure are called resonance structures. We can account for both structures by writing the resonance form, as shown below. All resonance structures have a double bond.

  27. Expansion into the d-orbital • Nonmetals with available d-orbitals (n> 3) can use these orbitals to hold more than 8 electrons. Unless Carbon is present, assume that these atoms are the central atoms. • When drawing Lewis structures of molecules with expanded octets, put octets around the outer atoms first, then account for any remaining electrons on the central atom. • Ex. SF4 F • After completing octets on the F atoms, there is a lone pair remaining on S 10 electrons around Sulfur F S F F

  28. Group Examples • Draw the Lewis structures for the following “expanded octet” molecules • PCl5 • SF6 • IF5 Always be mindful when dealing with atoms of n>3 of the possibility of surpassing 8 valence electrons. • Draw the resonance structures of: • ClO4- • PO43-

  29. Revisiting Dipoles: Polar vs. Nonpolar Molecules • Dipole moments describe the coulombic attraction that is formed between the partial positive and negative charges on nonmetals when electrons are unevenly shared. • Dipole moments are vectors, quantities with both magnitude and direction. Imagine pulling on a rope. The force is in the direction of the pull. • If two people pull on opposing ends of a rope with equal force, the net forceis zero. • Polar molecules have an overall dipole moment, nonpolar molecules don’t.

  30. Table of Electronegativities Electronegativity of atoms increases up and to the right.

  31. Here, electronegativity is labeled as β - + O O C C S O - δ δ δ Equal force of attraction in opposite directions: Non polar + = 0 Δβ = 0.89 Δβ = 0.89 Unequal dipole moments. Thus, there is an overal dipole: Polar - - + δ δ δ + = Δβ = 0.03 Δβ = 0.89

  32. Group Example • Draw a water molecule. Then draw the dipoles and determine the net dipole. Is the molecule polar? O H H + = net dipole = polar

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