1. Covalent Bonding Unit 6 Chapter 6

2. Haves and Have-Nots In the late 1800’s and early 1900’s, scientists did not know very much about the reasons that certain chemicals reacted with others. Knowledge of this reactivity, however, was shown to help in industrial and medicinal applications, but there was no real understanding, only a repertoire of techniques.

3. Ionic Bonding is Like Magnets We knew about electrons and could determine that cations lost electrons and anions gained electrons. We understood why ions were attracted to each other because we had physical theories of magnetism that could be used as models.

4. A Wrench Thrown in the Gears Positive and negative charges attracted each other – this gave us ionic bonding. Non-ionic bonding, when a nonmetal bonds with another nonmetal, was very different. We had very little understanding of this phenomenon.

5. Gilbert the Octopus Gilbert Newton Lewis studied the behavior of many ionic and non- ionic compounds and Came up with the Octet Rule in 1902. He envisioned atoms as cubes. Gilbert Newton Lewis (1875 – 1946)

6. Thinking Cubed In Lewis’ model, electrons were stationary & formed the corners of the cubes. A graduate student saw Lewis’ work & Suggested that the corners were shared This led to bonding. Lewis’ Original Diagrams

7. The Dot Before the Dot Com In 1916, Lewis published his paper on “The Atom and the Molecule” which discussed non-ionic bonding. In his paper, Lewis discussed ways to predict non-ionic bonding using diagrams. These diagrams would later become known as Lewis Dot Structures.

8. Languish and Langmuir Unfortunately, Lewis was a poor communicator. His ideas did not get very far. During WWI, he met Irving Langmuir Who built on Lewis’ work and published a paper in 1921 called “The Arrangement of Electrons in Atoms and Molecules.” Irving Langmuir (1881 – 1957)

9. I dunno Lewis…I’ll call it “Mine!” In his paper, Langmuir coined the term “covalence” to describe the sharing of electrons in non- ionic bonding. He promoted the Octet Rule and Covalent Bonding so well that the theory was often known as the Lewis-Langmuir theory (or simply Langmuir’s).

10. Lewis Dot Structures We can imagine an atom like a square. According to the Octet Rule, an atom is stable when it has 8 electrons surrounding it. Since there are four sides of a square, each side can hold 2 electrons. Ne

11. Lewis Dot Structures Continued Number of dots surrounding an element is determined by the # of valence electrons. Follow Hund’s Rule (from Quantum Mechanics) and put a dot on each side of the Atomic symbol until we have to pair them up. H B C N O F Ne Be

12. What’s it all mean? Unpaired electrons are free to participate in bonding. Pairs of electrons do not participate in bonding and are called Lone Pairs. (Lone Pear)

13. Bonding with Valence Electrons Covalent means “with valence” Bonds form between unpaired valence electrons of adjacent atoms Atoms will only make as many bonds as there are unpaired electrons N HH H Already Paired Up Do not Bond!

14. Redox Redux An Oxidation State is the charge an atom would have if the bonds within a molecule were completely ionic. Many reactions are driven forward by a change in oxidation state. When an atom’s oxidation state is increased (made more positive), it is oxidized. When an atom’s oxidation state is decreased (made more negative), it is reduced.

15. IUPAC Rules! Rules for determining Oxidation State: 1. Atoms in their elemental state have an oxidation state of 0. 2. Any simple monatomic ion has an oxidation state equivalent to the charge of the ion. 3. Hydrogen is (almost) always +1 and oxygen is (almost) always -2. 4. The sum of the oxidation states must equal the charge of the molecule/ion (0 in a neutral atom) A Charge is a Type Of Oxidation State!

16. In Practice A monatomic ion’s charge is its oxidation state. Hydrogen is +1 unless it is bonded to an active metal e.g. LiH, H = -1 Oxygen is -2 unless it is bonded to itself e.g. peroxides, H 2 O 2, O = -1 Representative elements typically acquire the same oxidation as if they were ions. e.g. Alkali Metals = +1, Halogens = -1 If there is uncertainty, the most electronegative element gets the negative charge e.g. FCl, F = -1, Cl = +1 This is NOT Common!

17. Oxidation Station The sum of the oxidation states must equal the charge of the molecule/ion. SCl 4 S = ?Cl = -1 Charge = 0 1 sulfur plus 4 chlorines = 0 S + 4(-1) = 0  S = +4 PO 4 -3 P = ?O = -2Charge = -3 P + 4(-2) = -3 P = +5

18. Make It So! Determine the oxidation state of sulfur in: SubstanceEquation Oxidation of S 1 S8S8S8S8 8(S) = 0 0 (elemental state) 2 H2SH2SH2SH2S 2(+1) + S = 0 -2 3 SO 2 S + 2(-2) = 0 +4 4 SO 3 -2 S + 3(-2) = -2 +4 5 H 2 SO 4 2(+1) + S + 4(-2) = 0 +6 A Charge is a type of Oxidation State, but An Oxidation State is not necessarily a Charge!