Electron Dot Structures (aka Lewis Diagrams)

1. Molecular Structure Electron Dot Structures (aka Lewis Diagrams)

2. Learning Objective • TLW express the arrangement of electrons in atoms through Lewis Valence Electron Dot Structures (TEKS 6.E.) • TLW be able to construct Lewis Dot Diagrams to illustrate ionic and compounds (TEKS 7.C)

3. Agenda • Background on Electron Dot Structures (aka Lewis Diagrams) • What they are • How they look • Octet Rule • Exceptions to the Octet Rule • Fruit Loop Activity - How to Draw Them • Resonance Structure • Bond Dissociation Energy • Group & Independent Practice

4. I. Electron Dot Structures – also known as Lewis Diagrams • Since valence electrons are the electrons involved in chemical bonding – they are significant • Chemists use a notation in which valence electrons are represented by dots and their bonds are represented by dashes • This notation known as Lewis Diagrams or Lewis Electron Dot Structures were named after their creator Gilbert Lewis

5. I. Electron Dot Structures A. Keeping track of valence electrons 1. How to write them for single atom? 2. Write the element’s symbol (“X”). 3. Put one dot for each valence electron on all four sides 4. Don’t pair up until they have to (Remember Hund’s rule?) X

6. B. The Electron Dot diagram for Nitrogen 1. Nitrogen has 5 valence e- 2. First we write the symbol. N 3. Then add 1 e- to each side 4. Until they are forced to pair up.

7. Electron Dots For Cations Metals will have few valence electrons (usually 4 or less) Ca

8. Electron Dots For Anions • Nonmetals will have many valence electrons; usually 5 or more • They will gain electrons to fill outer shell. P P

9. Practice Drawing Dot Diagrams for Elements • Hydrogen • Barium • Oxygen • Titanium • Neon • Independent Practice Set - link

10. Now Let’s Talk About Chemical Bonding

11. Ionic Bonds • An ionic bond is a chemical bond formed by the electron attraction between positive and negative ions. • Ionic bonds are made when an electron from the valence shell of one atom is transferred to the valence shell of another atom. Ionic bonds occur between metals and non-metals. • The atom that lost an electron becomes a positive ion and the atom that gains the electron becomes a negative ion.

12. Ionic Bonds • In NaCl the sodium ion has one less electron than protons so it has a positive charge. • The chlorine ion has one more electron than protons so it has a negative charge. • Since positives are attracted to negatives the two ions are attracted to each other. • The atom that loses an electron becomes a cation which is positive, and the atom that gains an electron becomes an anion which is negative. • The nature of the ionic bonds facilitates the formation of ionic solids by attracting other charged atoms to form a solid. • The ions are arranged in a crystalline structure with each Na+ ion attracted to several Cl- ions and each Cl- ion attracted to several Na+ ions. • There are no NaCl molecules.

14. Covalent Bonds • Covalent Bonds are chemical bonds formed by the sharing of a pair of electrons between atoms. Occur between non-metals. The nuclei of two different atoms are attracting the same electrons. • Therefore, unlike ionic bonds where an electron is moved from one atom to another the electrons are shared. • The Octet Rule is a tendency of atoms in molecules to have eight electrons in their valence shells. (Two for helium atoms.) The Octet Rule is a general rule, but is not followed by all molecules.

15. Covalent Bonds • Multiple Bonds are sometimes found in molecules so that the molecules satisfy the octet rule. • A single bond (which was discussed earlier) is when a single pair of electrons is shared between the two atoms. • A double bond is when two pairs of electrons are shared between two atoms. • A triple bond is when three pairs of electrons are shared between two atoms. (Notice a trend?) • Double and triple bonds mostly occur when the elements Carbon(C), Nitrogen(N), Oxygen(O) and Sulfur(S) are involved. • An example of a molecule with double bonds is Carbon Dioxide (CO2). Notice that each element ends up with eight electrons around it.

16. Polar Covalent Bonds • In bonds between atoms of the same element the sharing of the electrons is equal between the two atoms. • When two atoms of different elements make a bond, the electrons will not usually be shared equally. • The electrons are pulled more toward the more electronegative element. Electronegativity is the measure of the ability of an atom in a molecule to draw bonding electrons to itself. In general, electronegativity increases from bottom to top and left to right on the periodic table. Fluorine is the most electronegative element since it has a tendency to pick up electrons easily and hold on to them strongly. An element like cesium has a low electronegativity. • The unequal sharing of electrons is called a polar covalent bond. • The definition of a polar covalent bond is a covalent bond in which the bonding electrons spend more time near one atom than the other.

17. Metallic Bonds • Metallic bonding occurs in metallic substances.  • Atoms of metals are held together in this structure by the sharing effect of the electrons amongst all of the atoms.  This forms a "sea" or a "cloud" of free electrons that floats around the surface of metals.  This cloud of electrons explains many physical properties of metals.  • In particular, the fact that metals are good conductors can be easily visualized as electric charges are carried from one end of a metal to the other.  • The fact that metals are shiny can also be explained by this theory.  The luster of metals is in fact due to light (photons) bouncing off the "cloud" of free electrons and creating a shimmering effect.

18. Ne Octet Rule a. Most atoms form bonds in order to obtain 8 valence e- b. Full energy level stability ~ Noble Gases

19. H O H Electron Pairs Lone Pair (LP) Bond Pair (BP)

20. H O H Electron Pairs Lone Pair (LP) Bond Pair (BP)

21. Practice Drawing Compounds • LiF • CaO • SO2

22. 4. Exceptions to the Octet Rule • Exceptions: • Hydrogen  2 valence e- • Groups 1,2,3 get 2,4,6 valence e- • Expanded octet  more than 8 valence e- (e.g. S, P, Xe) • Radicals  odd # of valence e-

23. Exceptions to the octet rule • Hydrogen will have a maximum of 2 valence e- • Hydrogen will never be in the middle H O H

24. F Be F F Exceptions to the octet rule • Groups 1, 2, 3 get 2, 4, 6 valence e-

25. Exceptions to the octet rule • Total number of lone pairs (LP) and bond pairs (BP) associated with an atom cannot exceed the number of valence shell orbitals (n2, where n = row on periodic table) • n = 1 (H) has 1 LP + BP • n = 2 (B, C, N, O, F) max. LP + BP = 4 • n = 3 (Al, Si, P, S, Cl) max. LP + BP = 9

26. N O Exceptions to the octet rule • Expanded octet  more than 8 valence e- (e.g. Al, Si, S, P, Cl) Very unstable!!

27. F F F S F F F Exceptions to the octet rule • Radicals  odd # of valence e-

28. Drawing Lewis Diagrams for Compounds 1. Find total # of valence e- for all elements (remember to multiply by subscripts) 2. Arrange atoms - singular atom is usually in the middle (except H). 3. Form bonds between atoms 2 e- 4. Distribute remaining e- to give each atom an octet 5. If there aren’t enough e- to go around, form double or triple bonds. 6. There will not be any lone electrons

29. Other Tips • Hydrogen and halogens bond once • Family of Oxygen (Group 16) bonds twice • Family of Nitrogen (Group 15) bonds three times. As does Boron. • Family of Carbon (Group 14) bonds four times. More easily forms double or triple bonds. • Bond all atoms together first with single bonds, then add multiple bonds until the above rules are followed – or you run out of valence electrons

30. Still More Tips • HONC, pronounced “honk.” This is the way to remember that all Hydrogens have one and only one bond to them. Most Oxygens have two bonds to them. Most Nitrogens have three bonds to them, and most Carbons have four bonds to them. SO REMEMBER:       HONC 1, 2, 3, 4. • Carbon is always a central atom, except in diatomic molecules like carbon monoxide. • Hydrogen is always a terminal atom with only one bond and no dots. • Fluorine is also a terminal atom • The lowest electronegativity atom (NOT the closest to fluorine) is usually the central atom. • The structures are usually balanced around the central atom.

31. Another Mnemonic • CNOS – think of “See” My Nose • You have two nostrils • These four elements (carbon, nitrogen, oxygen, and sulfur) readily form double bonds using lone electrons

32. 6. Examples • CF4 1 C × 4e- = 4e- 4 F × 7e- = 28e- 32e- F F C F F - 8e- 24e-

33. BeCl2 1 Be × 2e- = 2e- 2 Cl × 7e- = 14e- 16e- ClBeCl - 4e- 12e-

34. CO2 1 C × 4e- = 4e- 2 O × 6e- = 12e- 16e- OCO - 4e- 12e-

35. Activity • Loopy Lewis Dot Diagrams • http://serc.carleton.edu/sp/mnstep/activities/19777.html • Instructions, Textbook References, Questions link

36. II. Polyatomic Ions A. To find total # of valence e-: 1. Add 1e- for each negative charge. 2. Subtract 1e- for each positive charge. B. Place brackets around the ion and label the charge.

37. C. Examples of polyatomic ions • ClO4- 1 Cl × 7e- = 7e- 4 O × 6e- = 24e- 31e- O O Cl O O + 1e- 32e- - 8e- 24e-

38. NH4+ 1 N × 5e- = 5e- 4 H × 1e- = 4e- 9e- H H N H H - 1e- 8e- - 8e- 0e-

39. III. Resonance Structure A. When more than one valid dot diagram is possible. 1. Consider the two ways to draw ozone (O3) 2. Which one is it? 3. Does it go back and forth? 4. It is a hybrid of both, like a mule; shown by a double-headed arrow

40. O O S O O O S O O O S O B. Example • SO3

41. IV. Bond Dissociation Energies • In covalent molecules, the compound molecule is more stable than the reactants • Total energy required to break the bond between to covalently bonded atoms is known as bond dissociation energy (BDE). • BDE is measured in kJ (kilo Joules) • BDE ranges from the weakest O – O at 142 kJ to strongest C = O at 1074 kJ)

42. Group Practice • Lewis Diagram Race… A Team Competition

43. Individual Practice • Lewis Dot Diagram – Practice 1 (link) • Yet Another Practice Set on Electron Dot Structure (Lewis Diagrams)