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Bellwork Tuesday

Bellwork Tuesday. List facts that you remember about ionic and covalent compounds. Unit 11 Ch7 & 8. Chapter 7 Ionic Bonding. Electron Configuration in Ionic Bonding. Valence Electrons – the electrons in the highest occupied energy level of an element’s atoms

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Bellwork Tuesday

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  1. Bellwork Tuesday List facts that you remember about ionic and covalent compounds

  2. Unit 11 Ch7 & 8 Chapter 7 Ionic Bonding

  3. Electron Configuration in Ionic Bonding • Valence Electrons – the electrons in the highest occupied energy level of an element’s atoms • To find the number of valence electrons look at the element’s group number 1 8 2 3 4 5 6 7

  4. Electron Configuration in Ionic Bonding • Electron dot structure - System of arranging dots representing valence electrons. • G.N. Lewis developed this system and so they are also referred to as Lewis Dot Structures.

  5. Example: • Write the electron configuration for phosphorous. 1s22s22p63s23p3 • Write the orbital diagram for phosphorous. ___ ___ ___ ___ ___ ___ ___ ___ ___ 1s 2s 2p 3s 3p • Write the Lewis dot structure for phosphorous.

  6. PracticeWrite Lewis dot structures for the following.

  7. Octet Rule • Octet rule – in forming compounds, elements tend to achieve the electron configuration of a noble gas. • Metals will typically lose electrons and therefore become positive Na ___ ___ ___ ___ ___ ___ 1s 2s 2p 3s Ne ___ ___ ___ ___ ___ ___ 1s 2s 2p 3s +

  8. Octet Rule • Nonmetals will typically gain electrons and therefore become negative O ___ ___ ___ ___ ___ 1s 2s 2p Ne ___ ___ ___ ___ ___ 1s 2s 2p 2-

  9. Electron Configuration in Ionic Bonding • Electron dot symbols, Lewis dot structures, can be used to represent ions. • EXAMPLES: • Na Cl F Mg O Al 2+ - - 2- 3+ +

  10. Chemical Valentines Dichromates are yellow, Copper solutions are blue, Make me a chemical valentine, And I will give points to you! You may make 2 Valentines (10 points each) due Monday Must be extra corny and super pretty.

  11. BellworkWednesday • On your way in, you should have received a card with an ion on it. Using your card, do the following… • Find a person or people who balance out the charge on your card and stand by them. • Write the compound you form with your group. • Name the compound. • Is your compound ionic or molecular? How do you know?

  12. BellworkMonday • Write out the Lewis dot structures for the following: • Ga • S • Br+3 • C-4

  13. 7.2 Ionic Bonds • Ionic Bond - A chemical bond formed by the electrostatic attraction between a cation and an anion when there is a transfer of electrons. • Ionic compounds are generally a metal and a nonmetal bonded together.

  14. Example: • When potassium reacts with chlorine what kind of compound is formed? • Ionic • Is there a transfer of electrons? • Yes • Using the Lewis structures show what happens to K and Cl when they combine to form salt.

  15. Properties of Ionic Compounds • High melting points • Solids at room temperature • Crystalline in structure, referred to as a crystal lattice MP: 802°C

  16. 7.3 Bonding in Metals • Metallic bonds – consist of the attraction of free-floating valence electrons for the positively charged metal ions. • These forces hold metals together Free Floating valence e- Positively charged metal ion

  17. Properties of Metallic Compounds: • Good conductors • Ductile – can be drawn into a wire • Malleable – can be hammered or forced into shapes

  18. Crystalline Structure of Metals • Metals and ionic compounds are in compact orderly patterns • Identically sized spheres have several arrangements • Simple cubic • Body-centered cubic • Face-centered cubic

  19. Bonding in Metals • Alloys – mixtures of two of more elements, at least two of which are metals. • Steels are the most important alloys today • Consist of Fe, C, B, Cr, Mn, Mo, Ni, W, V • Ex: Brass, steel, 14 K gold, sterling silver, & cast iron

  20. BellworkFriday • Draw Lewis structures for atoms of magnesium and sulfur. • Show how these atoms could combine to form a compound using the Lewis structures you drew.

  21. Unit 11 Chapter 8 Covalent Bonding

  22. 8.1 Covalent Bonds • Covalent bond– Occurs when a pair of electrons is shared between two atoms. • Often between two nonmetals. • Single covalent bond – two atoms share a single pair of electrons • Double covalent bond – two atoms share two pairs of electrons • Triple covalent bond – two atoms share three pairs of electrons H H O O N N

  23. Covalent Bonds • Lewis dot structures can be useful for representing covalent bonds between elements in a covalent compound. • H + H  • Cl + Cl  • O + O  • N + N  H H Cl Cl O O N N

  24. Rules for writing electron dot structures: (use pencil!!!) • 1. Add up the valence electrons from all the atoms in the compound. • Don’t try to keep track of which electrons come from which atoms. • If you are working with an ion, you must add or subtract electrons to account for the charge. • H2O (2)H + (1) O (2)(1e-) +(1) (6e-) = 8e-

  25. Rules for writing electron dot structures: • 2. Put the element that you have the fewest of as the central element. (Make it symmetrical) • Put the elements in spatial order. • H2O (2) H + (1) O H O H

  26. Rules for writing electron dot structures: • 3. Use a pair of electrons to form a bond between each pair of atoms. H O H

  27. Rules for writing electron dot structures: • 4. Arrange the remaining electrons to satisfy the duet rule for hydrogen and the octet rule for all remaining atoms. H O H

  28. Rules for writing electron dot structures: • 5. Count the number of electrons represented in the drawn molecule. • If two too many electrons are represented: • draw a double bond between two elements • remove a pair of electrons from each element taking place in the bond. H O H 1 2 3 4 5 6 7 8

  29. EXAMPLES: • CH4 • 1. (1) C + (4) H (1)(4e-) + (4)(1e-) = 8e- • 2. Spatial order • 3. Draw bonds • 4. Octet rule satisfied? • 5. # of e- match? H C H H H

  30. EXAMPLES: • CO2 • 1. (1) C + (2) O (1)(4e-) + (2)(6e-) = 16e- • 2. Spatial order • 3. Draw Bonds • 4. Octet rule satisfied? • 5. # of e- match? O C O

  31. EXAMPLES: • NH3 • 1. (1) N + (3)H (1)(5e-) + (3)(1e-) = 8e- • 2. Spatial order • 3. Draw bonds • 4. Octet rule satisfied? • 5. # of e- match? N H H H

  32. EXAMPLES: • CCl4 • 1. (1) C + (4) Cl (1)(4e-) + (4)(7e-) = 32e- • 2. Spatial Order • 3. Draw bonds • 4. Octet rule satisfied? • 5. # of e- match? Cl Cl Cl C Cl

  33. EXAMPLES: • NH4+ • 1. (1) N + (4) H - (1)(+) (1)(5e-)+ (4)(1e-) - (1)(1e-) = 8e- • 2. Spatial order • 3. Draw bonds • 4. Octet rule satisfied? • 5. # of e- match? H [ ]+ N H H H

  34. EXAMPLES: • SO42- • 1. (1) S + (4) O + (2)(-) (1)(6e-)+ (4)(6e-) + (2)(1e-) = 32e- • 2. Spatial Order • 3. Draw bonds • 4. Octet rule satisfied? • 5. # of e- match? O [ ] 2- O S O O

  35. EXAMPLES: • CN- • 1. (1) C + (1) N + (1)(-) (1)(4e-) + (1)(5e-)+ (1)(1e-) = 10e- • 2. Spatial order • 3. Draw Bonds • 4. Octet rule satisfied? • 5. # of e- match? [ ]- C N

  36. EXAMPLES: • CO32- • 1. (1) C + (3) O + (2)(-) (1)(4e-)+ (3)(6e-) + (2)(1e-) = 24e- • 2. Spatial Order • 3. Draw bonds • 4. Octet rule satisfied? • 5. # of e- match? [ ] 2- O C O O

  37. Exceptions to the Octet Rule:Elements that can have extra electrons! B Boron Trifluoride 3 bonds on B(6e-) F F F F F F S F F F Cl Phosphorous Pentachloride 5 bonds on P (10 e-) Sulfur Hexafluoride 6 bonds on S (12 e-) Cl P Cl Cl Cl

  38. REMEMBER: “P B S” bonded to ANY halogen can be an exception!

  39. Resonance Structures • Structures that can occur when it is possible to write two or more valid electron dot structures that satisfy the octet rule. EXAMPLES: • CO32- [ ] [ ] [ ] 2- 2- 2- O O O C C C O O O O O O

  40. Resonance Structures EXAMPLE: • NO3- [ ] [ ] - - [ ] - O O O N N N O O O O O O

  41. BellworkFriday • Draw the following Lewis dot structures. • CCl4 NH4+ • SO42- CN- • CO32-

  42. VSEPR • Valence Shell Electron Pair Repulsion theory • Valence electrons on the central atom repel each other.

  43. VSEPR: • Regions of electron density (where pairs of electrons are found) can be used to determine the shape of the molecule. • CO2 • Here there are two regions of electron density. • The regions want to be as far apart as possible, so it is linear. O C O

  44. EXAMPLES: 1 4 H • CH4 • There are four electron pairs. • You would expect that the bond angles would be 90° but… • Because the molecule is three-dimensional, the angles are 109.5°. • The molecule is of tetrahedralarrangement. C H H 2 H 3

  45. EXAMPLES: 1 4 • NH3 • Four regions of electron density • But one of the electron pairs is a lone pair • The shape is called trigonal pyramidal N H H 2 H 3

  46. EXAMPLES: 1 4 • H2O • Four regions of electron density • But two are lone pairs • This structure is referred to as bent O H H 2 3

  47. EXAMPLES: 1 3 [ ] 2- • CO32- • Three regions of electron density • This structure is referred to as trigonal planar C O O 2 O

  48. Practice determining molecular shape: • H2S • 4 regions of e- density • 2 lone pairs • bent S H S H H H

  49. Practice determining molecular shape: • SO2 • 3 areas of e- density • 1 lone pair • bent S O S O O O

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