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Covalent Bonding

Covalent Bonding. Unit 7 Chapter 6 & 7. 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.

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Covalent Bonding

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  1. Covalent Bonding Unit 7 Chapter 6 & 7

  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 N F B C Ne O 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 Already Paired Up Do not Bond! H H N H

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

  16. This is NOT Common! 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, H2O2, 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

  17. Oxidation Station • The sum of the oxidation states must equal the charge of the molecule/ion. • H2O • H = +1 O = -2 Charge = 0 • 2 hydrogens at +1 plus 1 oxygen at -2 = 0 • 2(+1) + 1(-2) = 0 • CCl4 • C = ? Cl = -1 Charge = 0 • C + 4(-1) = 0 • C = +4

  18. Make It So! Determine the oxidation state of sulfur in: A Charge is a type of Oxidation State, but An Oxidation State is not necessarily a Charge!

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