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UNIT 3

UNIT 3. Introduction to Lewis Structures and Covalent Bonding. Types of Chemical Bonds. Ionic bonds hold ions together primarily by electrostatic forces. Covalent bonds hold atoms together primarily by the sharing of valence electrons.

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UNIT 3

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  1. UNIT 3 Introduction to Lewis Structures and Covalent Bonding

  2. Types of Chemical Bonds • Ionic bondshold ions together primarily by electrostatic forces. • Covalent bondshold atoms together primarily by the sharing of valence electrons. • Metallic bonds.Here the atoms form a crystal and the valence electrons are free to move throughout the crystal. • We will now study covalent bonding.

  3. Covalent Bonds • Covalent bonds hold atoms together primarily by the sharing of valence electrons. • Valence electrons are electrons in the outer shell of an atom. • In general, the valence electrons are the electrons in the s and p orbitals of the highest shell.However, sometimes electrons in d orbitals also participate in bonding. • Example: Pb is [Xe]4f145d106s26p2 Note the filled 4f and 5d subshells. Pb has four valence electrons.

  4. Valence Electrons • For the main group elements, the number of valence electrons is the last digit of the group number. • Knowing the number of valence electrons allows us to draw Lewis symbols for the elements and Lewis structures for compounds.

  5. Lewis Symbols for Atoms and Monatomic Ions • Lewis symbols are very helpful in studying bonding because they show only the valence electrons of the atoms or ions. • A Lewis symbol consists of the symbol of the element or ion (include the charge of the ion) and one dot for each valence electron. • [Ne]3s1 [Ne]3s23p1 [Ne]3s23p3 [Ne]3s23p5 • Lewis symbols are mainly used for main group elements. Obviously, drawing Lewis symbols for transition elements is trickier.

  6. The Octet Rule Atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons (the s and p subshells are full). metals lose e- Na+ Mg2+ Al3+ nonmetals gain or lose e-:

  7. Covalent Bonding in the H2 Molecule The atoms in H2 are held together mainly because the two nuclei are electrostatically attracted to the concentration of negative charge between them. H∙ + H∙  H׃H -or- H∙ + H∙  H-H

  8. Lewis Structures of Covalent Compounds Knowing the number of valence electrons for an element and using the octet rule will let you successfully predict the bonding in many covalent compounds. Here are the rules to follow: 1. Decide which atoms are bonded. 2. Count all valence electrons. 3. Put two electrons in each bond. 4. Complete the octets of the atoms attached to the central atom except for H, which takes a duet. 5. Put any remaining electrons on the central atom. 6. If the central atom has less than an octet, form double or triple bonds.

  9. Lewis Structures of Covalent Compounds 1. Decide which atoms are bonded. 2. Count all valence electrons. 3. Put two electrons in each bond. 4. Complete the octets of the atoms attached to the central atom except for H, which takes a duet. 5. Put any remaining electrons on the central atom. 6. If the central atom has less than an octet, form double or triple bonds.

  10. Lewis Structures of Multiple Bonds 1. Decide which atoms are bonded. 2. Count all valence electrons. 3. Put two electrons in each bond. 4. Complete the octets of the atoms attached to the central atom except H, which takes a duet. 5. Put any remaining electrons on the central atom. 6. If the central atom has less than an octet, form double or triple bonds. Triple bonds are the shortest and strongest.

  11. Lewis Structures of Covalent Compounds 1. Decide which atoms are bonded. 2. Count all valence electrons. 3. Put two electrons in each bond. 4. Complete the octets of the atoms attached to the central atom, except H, which takes a duet. 5. Put any remaining electrons on the central atom. 6. If the central atom has less than an octet, form double or triple bonds.

  12. Electronegativity and Bond Polarity Ionic and covalent bonds are the extremes:complete control of the valence electrons and complete sharing of the valence electrons. Most bonds are somewhere in between. Electronegativityis the ability of an atom in a bondto attract electrons. Atoms that are more electronegative will tend to have a partial negative charge. In a molecule, a partial separation of charge means the covalent bond is polar. Atoms with high electron affinity and high ionization energy will be the most electronegative.

  13. Electronegativities can be found in your text. If the difference in electronegativities of the two atoms is between 0.4 and 2.0, the bond is polar (aka polar covalent). If the difference in electronegativities of the two atoms is ≤0.4, the bond is nonpolar.

  14. Electronegativity and Bond Polarity δ+ δ- ΔE.N. = 4.0 – 4.0 = 0.0 The bond is 100% covalent. The molecule is nonpolar. ΔE.N. = 4.0 – 2.1 = 1.9 The bond is polar covalent. The molecule is polar. ΔE.N. = 4.0 – 1.0 = 3.0 The bond is ionic. The larger the difference in electronegativity between two atoms, the more polar their bond. ΔE.N. = 2.0 is sometimes used as the “boundary” between ionic and polar covalent.

  15. Using Formal Charges to Decide Between Lewis Structures • The formal charge is the charge that an atom in a molecule would have if all atoms had the same electronegativity. • Formal charges are calculated after the Lewis structures are drawn. • Formal charge = # valence e- in atom - # bonds to atom - # unshared e- in atom • Lewis structures with the formal charges closest to zero and with the negative charges on the more electronegative atoms will be more stable. F.C. on H = 1 – 1 – 0 = 0 F.C. on O’s bonded to H’s = 6 – 2 – 4 = 0 F.C. on other O = 6 – 1 – 6 = -1 F.C. on P = 5 – 4 – 0 = +1

  16. Using Formal Charges to Decide Between Lewis Structures Formal charge = # valence e- in atom - # bonds to atom - # unshared e- in atom F.C. on H = 1 – 1 – 0 = 0 F.C. on O’s bonded to H’s = 6 – 2 – 4 = 0 F.C. on other O = 6 – 1 – 6 = -1 F.C. on P = 5 – 4 – 0 = +1 F.C. on H = 0 F.C. on O’s bonded to H’s = 0 F.C. on other O = 6 – 2 – 4 = 0 F.C. on P = 5 – 5 – 0 = 0 preferred structure

  17. Using Formal Charges to Decide Between Lewis Structures Formal charge = # valence e- in atom - # bonds to atom - # unshared e- in atom F.C. on H = 1 – 1 – 0 = 0 F.C. on O’s bonded to H’s = 6 – 2 – 4 = 0 F.C. on other O’s = 6 – 1 – 6 = -1 F.C. on S = 6 – 4 – 0 = +2 F.C. on H = 0 F.C. on O’s bonded to H’s = 0 F.C. on other O = 6 – 2 – 4 = 0 F.C. on S = 6 – 6 – 0 = 0 This second structure is the preferred one. It violates the octet rule, but agrees with experimental evidence, which shows two of the S-O bonds being shorter. Double bonds are shorter than single bonds. Elements in the third period have d orbitals available that can accommodate extra electrons.

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