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Lewis Symbols

Lewis Symbols. To help us to focus on the valence electrons – those that can participate in bonding - we use Lewis Symbols (in honor of scientist G.N. Lewis). Lewis Dot Symbols. Lewis Dot symbol (or Electron dot symbol) Dots placed around an element ’ s symbol represent valence electrons

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Lewis Symbols

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  1. Lewis Symbols • To help us to focus on the valence electrons – those that can participate in bonding - we use Lewis Symbols (in honor of scientist G.N. Lewis)

  2. Lewis Dot Symbols • Lewis Dot symbol (or Electron dot symbol) • Dots placed around an element’s symbol represent valence electrons • Pair electrons as needed • Ions are placed in brackets with charge outside • Easily shows “Octet rule” • Tendency of an atom to achieve an electron configuration having 8 valence electrons • Same as the electron configuration of a noble gas • The 8 electrons shown as 4 pairs

  3. Group Practice: • Write out dot diagrams showing the valence electrons of the following atoms. Use principles of electron configuration to predict and explain the ionic compoundseach pair will react to form: • Magnesium fluoride • Aluminum oxide • Nickel(II) chloride

  4. Molecular Compounds • Write the orbital diagram for a hydrogen atom. • We have seen evidence from gaseous reactions that elemental hydrogen exists as a diatomic molecule. Use your orbital diagram to explain why hydrogen atoms would be more stable (lower chemical energy) as H2 molecules.

  5. Electron Dot Diagrams for theDiatomic Elements • How does hydrogen obtain a noble-gas electron configuration?

  6. Covalent Compounds and Bonding • When two nonmetals form a compound, the bond between atoms is covalent. • Both atoms are close to the noble-gas electron configuration, so sharing electrons will allow both to obtain it. • In a covalent bond, each shared electron is attracted simultaneously to two nuclei.

  7. Covalent Bonds • Electrons not transferred in this case • Electrons typically shared in pairs

  8. Carbon Dioxide Example • The atoms of CO2 molecules are held together by strong covalent bonds. • No bonds connect the molecules, so CO2 molecules separate easily from each other into the gas state at room temperature. Figure 8.13 Figure from p. 28

  9. Whiteboard Practice • Use valence electron concepts and electron dot diagrams to represent these compounds: • HF • CF4 • O2 • CO2

  10. The Octet Rule • Just as in ionic bonding, covalent bonds are formed so that each atom can have the noble-gas electron configuration. Noble gases have 8 valence electrons, an octet.

  11. The Halogens • Do the atoms in each of these molecules have an octet? • Why do the halogens exist as diatomic molecules?

  12. Multiple Bonds • How many valence electrons does an oxygen atom have? • How many does it need to obtain an octet? • O2 has a double bond, two pairs of shared electrons • How many valence electrons does a nitrogen atom have? • How many does it need to obtain an octet? • N2 has a triple bond, three pairs of shared electrons

  13. The Octet • An unreactive or stable compound usually has the maximum number of valence electrons per core (8) • Same as the electron configuration of a noble gas • Covalently bonded atoms achieve 8 valence electrons by sharing electrons • The 8 electrons exist in 4 pairs • H atoms bond with other atoms to obtain a total of 2 electrons like He (duet).

  14. Practice: Valence Electrons and Number of Bonds • How many bonds do each of the following atoms tend to form? • H • Cl • O • N • C

  15. Carbon Compounds Figure 8.21 • Carbon has: • Four valence electrons • The ability to form four bonds • The ability to bond to itself • Very strong bonds when bonded to itself • Carbon molecules are ubiquitous in nature. • Aside from what I’ve mentioned here, skip “Bonding in Carbon Compounds”, pp. 307-310 (top).

  16. Hydrocarbons • Aromatic hydrocarbons • A class of hydrocarbons which has carbon atoms arranged in a six-atom ring with alternating single and double bonds • Delocalized structures Figure 8.22 Figure 8.22

  17. Functional Groups in Hydrocarbons

  18. Ionic and Covalent • In ionic compounds, ions are held together by electrostatic forces – forces between oppositely charged ions. • In molecular compounds, atoms are held together by covalent bonds in which electrons are shared. Figure 8.2

  19. Activity: Identifying Types of Bonding • Identify the type of bonding in each of the following substances: • NaF • ClO2 • FeSO4 • SO2 • Ca(ClO2)2

  20. Steps for Writing Lewis Structures • Write an atomic skeleton. • Sum the valence electrons from each atom to get the total number of valence electrons. • Place two electrons, a single bond, between each pair of bonded atoms (can also be drawn as a line) • Place remaining valence electrons to complete the octet of each outer atom. These are called non-bonding electrons or lone pairs. If there are “extra” electrons, place the pairs around the central atom. • If necessary to satisfy the octet rule, shift unshared electrons from non-bonded positions on atoms with completed octets to positions between atoms to make double or triple bonds.

  21. Activity: Lewis Structures • Draw Lewis structures to show how electrons are shared in these molecules. • C2H6 • C2H4 • C2H2 • HCN • CO2 • NH3

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