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Covalent Bonding & Molecular Compounds

Covalent Bonding & Molecular Compounds. Depicting Molecular Compounds. Key Terms. A molecule is a neutral group of atoms that are held together by covalent bonds. A molecular compound is a chemical compound whose simplest units are molecules.

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Covalent Bonding & Molecular Compounds

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  1. Covalent Bonding & Molecular Compounds Depicting Molecular Compounds

  2. Key Terms • A moleculeis a neutral group of atoms that are held together by covalent bonds. • A molecular compoundis a chemical compound whose simplest units are molecules. • A chemical formulaindicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts. • A molecular formulashows the types and numbers of atoms combined in a single molecule of a molecular compound. • A diatomic moleculeis a molecule containing only two atoms.

  3. Formation of a Covalent Bond • Describe what happens to potential energy as a bond forms. • As a covalent bond forms, the attractive forces between the atoms overpower the repulsive forces between the atoms. • When the bond has formed, potential energy is at a minimum.

  4. Formation of a Covalent Bond

  5. Characteristics of Covalent Bonds • Bond lengthis the average distance between two bonded atoms. • By sharing electrons, each H atom has a full 1s orbital. • Bond energyis the energy required to break a chemical bond and form neutral, isolated atoms.

  6. Bond Length vs. Bond Energy • As bond length decreases, bond energy increases. • As bond length increases, bond energy decreases. • Therefore, the relationship between bond length and bond energy is an inverse relationship.

  7. The Octet Rule • Chemical compounds tend to form so that each atom has an octet of electrons in its highest occupied energy level. • Octet = 8 electrons. • An octet results in fully filled s + p orbitals.

  8. Exceptions to the Octet Rule • 1. Some elements are satisfied with fewer than 8 valence electrons: • Hydrogen only needs 2 valence electrons (1 bond). • Beryllium only needs 4 valence electrons (2 bonds). • Boron only needs 6 valence electrons (3 bonds). • 2. Some elements can be surrounded by more than eight electrons when they bind to highly electronegative atoms (F, O, Cl). Called an expanded octet.

  9. Lewis Structures • Electron-dot notation used to represent molecules. • Dotsrepresent unshared electrons. : • Dashesrepresent shared electron pairs. - • A shared pair of dots (electrons) is replaced with a dash. • A single dash represents a single bond.

  10. Steps for Drawing Lewis Structures • 1. Determine the type and number of atoms in the molecule. • For CH3I the atoms are: • 1 carbon, 3 hydrogen and 1 iodine. • Obtain the correct element cards from the bag.

  11. Steps for Drawing Lewis Structures • 2. Write the electron-dot notation for each type of atom and then determine the total number of valence electrons. • Use different colored beads to represent the valence electrons around each type of atom. • C = 1 atom x 4 valence electrons = 4 val e- • H = 3 atoms x 1 valence electron = 3 val e- • I = 1 atom x 7 valence electrons = 7 val e- • TOTAL = 14 valence electrons

  12. Steps for Drawing Lewis Structures • 3. Place atom with most open spaces (single electrons) in the middle. • Carbon will always be in the center if it is present. • Hydrogen and halogens will never be centrally located. • Arrange cards around the central atom.

  13. Steps for Drawing Lewis Structures • 4a. Make sure each element (except hydrogen) is surrounded by eight valence electrons. • If not all atoms are surrounded by 8 valence electrons, move electron pairs to form double or triple bonds. • Hydrogen and halogens will never form double or triple bonds.

  14. Steps for Drawing Lewis Structures • 4b. Replace bonding electron pairs with dashes. • Replace bonding electron pairs (beads) with straw sticks • 5. Check work by counting number of valence electrons. • Should have same number of electrons as in step 2.

  15. Double Bond Example CH2O • Step 1: • 1carbon, 2 hydrogen, 1 oxygen • Step 2: • 12 total valence electrons. • Step 3:

  16. Double Bond Example • Step 4a: • Carbon and oxygen are both one electron short. Move an electron pair between oxygen and carbon. • Step 4b: • Step 5: • Final structure has 12 valence electrons (same as in step 2).

  17. Triple Bond Example CO • Step 1: • 1carbon & 1 oxygen • Step 2: • 10 total valence electrons. • Step 3:

  18. Triple Bond Example • Step 4a: • Move two electron pairs between oxygen and carbon to satisfy octet rule for both. • Step 4b: • Step 5: • Final structure has 10 valence electrons (same as in step 2).

  19. Single vs. Multiple Bonds • Single bondsconsist of one shared electron pair. • Multiple bondsconsist of two of more shared electron pairs. • Double bonds have two shared electron pairs. • Have shorter bond length than single bonds. • Have greater bond energy than single bonds. • Triple bonds have three shared electron pairs. • Have shorter bond length than double bonds. • Have greater bond energy than double bonds.

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