Chemical Bonds, Electronegativity , Lewis Structures, VSEPR and Molecular Geometry

1. Chemical Bonds, Electronegativity, Lewis Structures, VSEPR and Molecular Geometry Students will identify differences between ionic and covalent bonding, draw lewis structures and identify molecular geometry using VSEPR.

2. Bonds • Force that holds groups of two or more atoms together and makes them function as a unit. • Bond Energy: energy required to break the bond

3. Ionic Bonding • Strong forces present from attractions among closely packed, oppositely charged ions. • Ionic Compounds: result when a metal reacts with a nonmetal

4. Covalent Bonds • Electrons are shared by nuclei. Bonding results from the mutual attraction of the two nuclei for the shared electrons.

5. Polar Covalent Bonds • Unequal sharing of electrons • Hydrogen fluoride is an example of this bond. • A delta δ is used to indicate a partial or fractional charge.

6. Bond Polarity • Has important chemical implications. • Used to assign a number that indicates an atom’s ability to attract shared electrons. We use electronegativity.

7. Electronegativity • Unequal sharing of electrons between 2 atoms. • The relative ability of an atom in a molecule to attract shared electrons to itself.

8. Bonds by Electronegativity • Zero: Covalent • Or Non Polar Covalent • Intermediate: Polar Covalent .5-1.7 • Large: Ionic 1.7+

9. Examples • Choose the bond that will be more polar, based on the electronegativity difference: • H—P , H—C • O—F, O—I • N—O , S—O • N—H , Si—H

10. Answers • H—C • O—I • S—O • N—H

11. Dipole Moment • Molecule with a positive center and a negative center. • Represented by an arrow which points towards the negative charged center. • Any diatomic molecule that has a polar bond has a dipole moment.

12. Lewis Structures • Representation of a molecule that shows how the valence electrons are arranged among the atoms in the molecule.

13. Steps for Writing Lewis Structures • Add up the valence electrons from all atoms. • Use one pair of electrons to form a bond between each pair of bound atoms. [use a line instead of dots] • Arrange the remaining electrons to satisfy the duet rule/octet rule.

14. Potassium Bromide Potassium loses its valence electron and bromine has 8 because is has a filled valence shell.

15. Duet Rule • Hydrogen forms stable molecules where it shares 2 electrons. For the HOBrINFCl’s. • H:H

16. Covalent Bonds • Water • Ammonia • Ammonium Ion

17. Practice – draw the Lewis dot structures for the following: • Methane • Hydrogen sulfide • Arsenic triiodide • Hydrogen selenide • Oxygen difluoride

18. Methane

19. Hydrogen sulfide

20. Arsenic triiodide

21. Hydrogen selenide

22. Oxygen difluoride

23. Multiple Bonds • Sum up the valence electrons. Ex. Carbon Dioxide • C [4] + O [6] + O [6] = 16 • Form a bond between the carbon and each oxygen. • O—C—O • Distribute the remaining electrons to achieve noble gas electron configurations. 16-4 = 12 • :Ö—C—Ö:

24. Was that correct? • NO.Didn’t follow the octet rule, so to achieve that, we must place double bonds between the atoms. • Ö—C—Ö • Can have double and triple bonds. • If a molecule can have more than one Lewis structure, it’s called resonance.

25. Molecular Structure • 3-dimensional arrangement of the atoms in a molecule. • Bent – “V-shaped” for a bond angle of 105° [water] • Linear – straight line – 180° [carbon dioxide] • Trigonal planar [flat] – all 4 atoms in the same plane with 120° bond angles. [also called triangular] • Tetrahedral [tetrahedron] – 4 identical triangular faces [methane – CH4]

26. VSEPR Model • Used to predict approximate structure of a molecule. • Valence, Shell, Electron, Pair, Repulsion Model

27. How? • The starting point is the Lewis dot structure of the compound. The number of electron pairs around the central atom determines the electron pair geometry. Lone pairs and bonding electron pairs are placed around the central atom. • The final step is to leave out lone pairs and determine the molecular geometry or molecular shape.

29. 1. Write the Lewis structure.2. Determine # of electron groups around the central atom.3. Determine the geometry. • NEG Electron Group Geometry • 2 linear • 3 trigonal planar • 4 tetrahedral • 5 trigonalbipyramidal • 6 octahedral

30. Rules • 1. 2 pairs of electrons on a central atom in a molecule are always placed 180° apart. Linear. • 2. 3 pairs of electrons on a central atom are always placed 120° apart. Trigonal. • 3. 4 pairs of electrons are always placed 109.5° apart. Tetrahedral. • See paige 389 in text

31. How about double bonds? • The same rules apply to double bonds.