2. Unit 12 Chemical Bonding & Molecular Geometry

3. Definitions Chemical Bonds Force that holds atoms together It’s all about the electrons (e-) Electrons are attracted to the nucleus of another atom

4. Types of Chemical Bonds Ionic Bond Bond between metal and nonmetal Attraction between positive cation and negative anion Electrons are transferred from metal to nonmetal

5. Types of Chemical Bonds Covalent Bonds Bonds in which e- are shared Most common type of bond

6. Types of Chemical Bonds Metallic Bonds Atoms are bonded to one another (not to other elements) Positive ions in a “sea” of negative charge (e-)

7. Definitions Octet rule (Rule of 8) 8 e- in the outer shell  very stable Exceptions: H 2 and He want a “duet”

8. Examples of Bonding Types Ionic Bonding: NaCl, K 2 S, Ca(NO 3 ) 2 Covalent Bonding H 2, Cl 2 Metallic Bonding Cu, Ag

9. Lewis Dot Diagrams A Lewis dot diagram depicts an atom as its symbol and its valence electrons. Ex: Carbon Carbon has four electrons in its valence shell (carbon is in group 14), so we place four dots representing those four valence electrons around the symbol for carbon.

10. Drawing Lewis Dot Diagrams Electrons are placed one at a time in a clockwise manner around the symbol in the north, east, south and west positions, only doubling up if there are five or more valence electrons. Same group # = Same Lewis Dot structure Ex. F, Cl, Br, I, At Example: Chlorine (7 valence electrons b/c it is in group 17)

11. Paired and Unpaired Electrons As we can see from the chlorine example, there are six electrons that are paired up and one that is unpaired. When it comes to bonding, atoms tend to pair up unpaired electrons. A bond that forms when one atom gives an unpaired electron to another atom is called an ionic bond. A bond that forms when atoms share unpaired electrons between each other is called a covalent bond.

12. Writing Lewis Dots Structures for Ions Uses either 0 or 8 dots, brackets and a superscript charge designate to ionic charge Ex.) Li +, Be +2, B +3, N -3, O -2, F -1

13. Writing Lewis Dots Structures (Ionic Compounds) Lewis Dot Diagrams of Ionic Compounds Ex. 1) NaCl Ex. 2) MgF 2

14. A substance made up of atoms which are held together by covalent bonds is a covalent compound. They are also called molecules. Lewis Dot Diagrams for Covalent Compounds

15. Covalent Compounds and Lewis Dot Diagrams Diagrams show bonds in a covalent compound and tells us how the atoms will combine Shared e - = bonding e - Non-shared e - = lone pair e - (a.k.a. non-bonding e - ) Ex. F 2

16. Drawing Electron Dot Diagrams for Molecules Chemists usually denote a shared pair of electrons as a straight line. F Sometimes the nonbonding pair of electrons are left off of the electron dot diagram for a molecule

17. Examples CH 4 CF 4

18. Types of Covalent Bonds Single Bond 2 e- are shared in a bond (1 from each atom) Double Bond 2 pairs of e- are shared (4 e- total, 2 from each atom) Triple Bond 3 pairs of e- are shared (6 e- total, 3 from each atom)

19. Rules for Drawing Lewis Dot Diagrams 1. Add up the total number of valence e - for each atom in the molecule. Each (-) sign counts as 1 e -, each (+) sign subtracts one e - 2. Write the symbol for the central atom then use one pair of e - to form bonds between the central atom and the remaining atoms. (see hints) 3. Count the number of e - remaining and distribute according to octet rule (or the “duet” rule for hydrogen) a) If there are not enough pairs, make sure the most electronegative elements are satisfied. Then, start shifting pairs into double and triple bonds to satisfy the octet rule. b) If there are extra e -, stick them on the central atom.

20. Hints: H is NEVER a central atom! Carbon will always be a central atom! Halogens (Group 17) are usually not central atoms. If you only have 1 of a certain element, it is usually the central atom.

21. Checking Your Work! But Remember.... The Structure MUST Have: the right number of atoms for each element, the right number of electrons, the right overall charge, and 8 electrons around each atom (ideally).

22. Covalent Compounds and Lewis Dot Diagrams HF NH 3 CCl 4 NH 4 + CO 2 NO 3 - *****

23. Resonance Structures Definition: When a single Lewis structure does not adequately represent a substance, the true structure is intermediate between two or more structures which are called resonance structures. Resonance Structures are created by moving electrons (in double or triple bonds), NOT atoms.

24. Resonance Structure Example, SO 2 Central atom = S This leads to the following structures: These equivalent structures are called RESONANCE STRUCTURES. The true structure is a HYBRID of the two. Arrow means “in resonance with”

25. Resonance Structure Example, NO 3 - Draw the Lewis diagram for NO 3 - with all possible resonance structures.

26. Molecular Geometry Molecular Geometry describes the 3-D arrangement of atoms in a molecule. We will use VSEPR theory to predict these 3-D shapes!

27. VSEPR: Shapes of Molecules VSEPR Theory (definition) “Valence Shell Electron Pair Repulsion” Based on idea that e- pairs want to be as far apart as possible Minimize electron pair repulsion  Gives molecule its shape

28. There is a fundamental geometry that corresponds to the total number of electron pairs around the central atom: 2, 3, 4, 5 and 6 lineartrigonal planar tetrahedraltrigonal bipyramidal octahedral

29. Basic Electron Pair Geometries Shapes Sum of Bonded Atoms & Lone e - 1. Linear 2. Trigonal planar 3. Tetrahedral

30. To determine the molecular geometry: 1. Draw the Lewis structure. 2. Count the number of bonded (X) atoms and non-bonded or lone pairs (E) around the central atom. 3. Based on the number of X and E, assign the molecular geometry. 4. Multiple bonds count as one bonded pair!

31. Molecular Geometry Notation A: Central Atom X: Bonded Atom E: Non-bonding electron pair (Lone pair e - on central atom)

32. Molecular Geometry: Two Bonded Items Total Bonds to C.A. # bonded atoms (X) # lone pairs (E) AXE Notation Molecular Geometry Bond Angles Example 220AX 2 Linear180°BeCl 2 Electron Pair Geometry = linear

33. Molecular Geometry: Three Bonded Items Total Bonds to C.A. # bonded atoms (X) # lone pairs (E) AXE Notation Molecular Geometry Bond Angles Example 3 30AX 3 Trigonal Planar 120°BCl 3 21AX 2 EBent118°NO 2 - Electron Pair Geometry = trigonal planar

34. Molecular Geometry: Four Bonded Items Total Bonds to C.A. Bonded atoms (X) Lone pairs (E) AXE Notation Molecular Geometry Bond Angles Example 4 40AX 4 Tetrahedral 109.5° CCl 4 31AX 3 E Trigonal pyramid 107° NH 3 22AX 2 E 2 Bent 104.5° H2OH2O e - pair geometry = tetrahedral

35. Molecules with More than One Central Atom Determine geometry for each central atom separately! Example: In acetic acid, CH 3 COOH, there are three central atoms:

36. Molecules with only two atoms are always linear! Examples: HClN 2

37. VSEPR Examples: What shape would the following compounds have according to VSEPR theory? CH 4 CO 2

38. O 2 CFCl 3 What shape (molecular geometry) would the following compounds have according to VSEPR theory?

39. H 2 S C 2 H 2 (hint: classify each C separately) What shape (molecular geometry) would the following compounds have according to VSEPR theory?

40. Bond and Molecule Polarity Polar Bond Covalent bond in which the electrons are unequally shared Ex. H 2 O Non-polar Bond Covalent bond in which the electrons are equally shared Ex. F 2 or CH 4 Predicting Bond Polarity Use Electronegativity!! (see next slide)

41. Predicting Bond Polarity Calculate the difference between the Pauling electronegativity values for the 2 elements Type of Bond IONIC (COVALENT) POLARNON-POLAR Types of Atoms 1 metal & 1 nonmetal (ex. NaCl) (generally) 2 nonmetals Ex. NH 3, H 2 O (generally) 2 nonmetals Ex. CCl 4, O 2 Electronegativity Difference ≥ 1.7> 0.4 but < 1.7≤ 0.4 0 – 0.4  Non-polar covalent 0.4 – 1.7  Polar covalent (more e/n element has greater pull) 1.7 and up  Ionic (e - are transferred between atoms)

42. Polar Molecules Polar Molecules (dipole) Molecule with separate centers of (+) and (-) charge In other words, molecules are polar if the pull in any one direction is not balanced out by an equal & opposite pull in the opposite direction

43. Polar Bonds and Polar Molecules Drawing Polar Molecules Positive and Negative regions shown by “delta”(δ) Ex. CH 3 Cl

44. Determining the Polarity of a Molecule Shape is crucial (determine the VSEPR shape 1 st ) All non-polar bonds = nonpolar molecule Polar bonds  see if they cancel each other out If they all cancel = nonpolar molecule If they are unbalanced = polar molecule

45. Determining Molecular Polarity Nonpolar Polar

46. Examples: Polar or non-polar? Determine if the following molecules are polar or nonpolar. H 2 S F 2 H 2 O

47. Special Types of Bonding Hydrogen Bonding Force in which a hydrogen atom covalently bonded to a highly electronegative element (F, O, or N) is simultaneously attracted to a neighboring nonmetal atom

48. Hydrogen Bonding Elements that undergo H-bonding Hydrogen bonding is FON! (Fluorine, Oxygen, and Nitrogen) Effects on Physical Properties H 2 O is most notable example of H-bonds Ice forms rigid, open structures  Increases volume upon freezing (floats) Molecules w/ higher molar mass have lower BP than H 2 O

49. Special Types of Bonding Van der Waals (London Dispersion) Forces Intermolecular force between the molecules of a substance Force of attraction between an instantaneous and induced dipole Molecules “make these up” (more or less)

50. Solids Classes of Solids Molecular Formed by molecules containing covalently bonded atoms Ionic Formed by cations and anions Network Covalent Formed by atoms, usually from Group IV A (Group 14) Metallic Formed by positive ions in a “sea” of electrons

51. Solids Comparison of Solids Type of Solid HardnessMalleabilityConductivity Melting Point Ease of Phase Change Molecular (2 nonmetals) SoftShattersInsulatorLow Easy to convert to gas Ionic (1 metal, 1 non) HardShatters Conducts if melted or in water High Difficult to convert to gas, solid Network Covalent Very HardShatters Poor Conductor High Difficult to convert to gas, solid Metallic Varies from soft to hard Very Malleable Good conductor as liquid or solid Usually High Difficult to convert to gas, Easy to covert to solid