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Unit 1: Structure and Properties of Matter

Chapter 4: Chemical Bonding. Unit 1: Structure and Properties of Matter. 4.1B Lewis Structures and Exceptions to Octet Rule. What We Did Last Time. What we did last time…. We learned about the duet and octet rules and how they determine the type of bonding that atoms participate in.

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Unit 1: Structure and Properties of Matter

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  1. Chapter 4: Chemical Bonding Unit 1: Structure and Properties of Matter 4.1B Lewis Structures and Exceptions to Octet Rule

  2. What We Did Last Time What we did last time…. • We learned about the duet and octet rules and how they determine the type of bonding that atoms participate in

  3. Learning Goals After this lesson, you will be able to: • Use the Lewis theory of bonding and the duet and octet rules to predict the bonding arrangements in different compounds • Explain exceptions to the octet rule

  4. Drawing Lewis Structures

  5. Drawing Lewis Structures

  6. Structure of a Molecule Example: Methanal, H2CO Step 1: The central atom is carbon, since it has the highest bonding capacity. It will be surrounded by the hydrogens and the oxygen. Step 2: The total number of valence electrons is: (2x1) + 6 + 4 = 12

  7. Structure of a Molecule Step 3: Now we place one pair of electrons between the central carbon and each attached atom (thereby using 6 of the 12 electrons). Step 4: Now we place lone pairs around each surrounding atom to obey the duet and octet rules. All we need are 6 electrons around the oxygen. We have now used all 12 electrons.

  8. Structure of a Molecule Step 5: There are no available electrons left since all 12 have already been placed. Step 6: No electrons can be placed on the central carbon since they have all been used. Step 7: The central carbon is short of valence electrons. It only has 6 and does not obey the octet rule. So we must move a pair of lone electrons from the oxygen and share that pair with the carbon (i.e. we create a double bond).

  9. Structure of a Molecule Step 8: There are no more electrons left to place in this case. Step 9: Now we replace the Lewis structure with the structural formula (with each covalent bond shown as a line). The lone pairs may or may not be shown.

  10. Structure of a Polyatomic Ion Example: Nitrate ion, NO3- Step 1: The central atom in an oxygen-rich ion is the nitrogen. It is surrounded by all of the oxygen atoms, which are directly attached to the nitrogen. Step 2: The total number of valence electrons is: 5 + 3(6) + 1 = 24 This includes the one additional electron due to the negative charge on the ion.

  11. Structure of a Polyatomic Ion Step 3: Now we place one pair of electrons between the central nitrogen and each attached atom (thereby using 6 of the 24 electrons). Step 4: Now we place lone pairs around each surrounding atom to obey the octet rule. We need 6 electrons around each oxygen. We have now used all 24 electrons.

  12. Structure of a Polyatomic Ion Step 5: There are no available electrons left since all 24 have already been placed. Step 6: No electrons can be placed on the central nitrogen since they have all been used. Step 7: The central nitrogen is short of valence electrons. It only has 6 and does not obey the octet rule. So we must move a pair of lone electrons from one of the oxygen atoms and share that pair with the nitrogen (i.e. we create a double bond).

  13. Exceptions to Octet Rule Most of the time, atoms obey the octet rule and we can write predictable structures for compounds. However there are exceptions in which the rule is not obeyed. There are two types of exceptions: • A central atom has an underfilled valence shell with fewer than 8 electrons • A central has an overfilled valence shell with more than 8 electrons

  14. Atoms with Underfilled Shells Atoms that commonly have underfilled valence shells include boron and beryllium. Ex. BF3 and BeF2 Boron and beryllium have only 2 or 3 valence electrons and cannot form more bonds. So they cannot fill their octets. However, these compounds are very reactive due to their deficiency of electrons. They readily react with compounds that are able to share lone electron pairs.

  15. Atoms with Overfilled Shells Some atoms have a tendency to exceed the octet rule, such as sulfur or phosphorus. Ex. SF6 and PCl5 Even though sulfur has only two vacancies in its 3p orbitals, it is able to form this high number of bonds by using its empty 3d orbitals. The orbital diagram of sulfur is shown here.

  16. Learning Check Answer the following questions to see if you met today’s learning goals. • Draw a Lewis structure for the compound CO2 and for the sulfate ion (SO42-). • Give examples of central atoms that often do not obey the octet rule. What kind of compounds do they form?

  17. Practice Exercises Complete the following exercises to make sure you understood what we did today. P. 205: #4, 5, 6

  18. Next Step Here is what comes next…. • Next time we will learn how the three-dimensional shapes of molecules are determined by electron structure.

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