1 / 19

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.

saburo
Télécharger la présentation

Covalent Bonding & Molecular Compounds

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  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.

More Related