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  1. Chapter 1 Prep - Chapter 1, page 7

  2. Chapter 1 Prep - Chapter 1, page 13 mercury (II) iodide aluminium oxide • 3. a) b) c) d) 4. a) b) c) d) 5. a) b) c) d) 6. a) b) c) d) 7. a) b) c) d) 8. a) b) c) d) sodium phosphide calcium bromide FeS ZnO MgI2 CoCl3 sulfur hexafluoride dinitrogen penta-oxide phosphorus pentachloride carbon tetrafluoride H2O never call it this! SO3 N2O4 N2O potassium phosphate ammonium chloride lithium perchlorate sodium hydrogen carbonate MgOOCCOO KClO NaCN (NH4)2SO4

  3. Chemistry 20 Chapter 1 PowerPoint presentation by R. Schultz robert.schultz@ei.educ.ab.ca

  4. 1.1 Forming and Representing Compounds • Chemical bonding theories developed by observing chemical reactions and physical properties of natural substances • Many patterns have been observed – one set, page 17 of Inquiry into Chemistry • You studied all of these in Science 10: • Metals found in combination with non-metals • Metals in pure form are solids ionic compounds

  5. 1.1 Forming and Representing Compounds • Non-metals combine to form solids, liquids, and gases • Only Noble Gases exist as individual atoms(we write metals as single atoms, but this is not how they exist) • Watch Chemical Families Video molecular compounds

  6. 1.1 Forming and Representing Compounds • Lewis Dot Diagrams: • show valence electrons only • valence electrons arranged into orbitals • Orbital: area of space where an electron of a given energy is most likely found - not what it is doing • What parts of atom does atomic symbol represent? (2 things) nucleus inner, non-valence electrons

  7. 1.1 Forming and Representing Compounds • Main group elementswill always have 4valence orbitals • Each orbital can hold maximum of 2 electrons • Electrons in completely filled orbital called lone pair • Electrons in half-filled orbital called bonding electrons Fig 1.5, p. 19

  8. 1.1 Forming and Representing Compounds • To draw a Lewis Dot Diagram of an element: • Write the element symbol • Count # of valence electrons • Each side of the symbol represents 1 of the 4 valence orbitals • Start on 1 of the sides, proceed around putting 1 electron on each, not doubling up until you have to

  9. bonding electrons lone pairs 1.1 Forming and Representing Compounds • Example: oxygen (6 valence electrons) • • Now you can start doubling up – electrons repel each other so 1 electron per orbital is more stable than 2 • • O • • It doesn’t matter which side lone pairs and bonding pairs are (it depends where you start counting from) important thing is that O ends up with 2 lone pairs and 2 bonding electrons

  10. • • • use sides not corners O • • 1.1 Forming and Representing Compounds • Please don’t do your Lewis Diagrams like this: • or like this: • • • electron pairs should be parallel to the symbol sides O • • •

  11. 1.1 Forming and Representing Compounds • Electronegativity – find on periodic table • Do BLM 1.1.3B – Lewis Diagrams of Atoms • Correct • Note that Noble Gases have the stable octet electron configuration

  12. 1.1 Forming and Representing Compounds • The stable octet configuration is what atoms attempt to achieve through bonding • Atoms can get the stable octet by transferring electrons or by sharing them

  13. to get the stable octet for each 1.1 Forming and Representing Compounds • Ionic compounds – transferring electrons • In reactions between metals and non-metals, • metals tend to lose electrons • non-metals tend to gain electrons We’ll discuss why later in the chapter

  14. • ‾ Na + Cl • • • • • • electron transfer If I ask you to do this on an assignment, quiz, or test, you must show everything circled in red 1.1 Forming and Representing Compounds • Example: • • + Na + Cl • • • • • • both have stable octet (Na has 8 electrons in the previous valence level) neither has stable octet

  15. • 2+ ‾ • • Ca F +2 • • • • On an assignment, quiz, or test, it is sufficient to do the following: • • • • • F • • • 2+ ‾ • + 2 • Ca + Ca F • • • • • • • • • • F • • • • 1.1 Forming and Representing Compounds • Example: + • • • • • • • • ‾ • + F • • F Ca + F • Ca + • • • • • • • • • • • • • F gets the stable octet, but not Ca, transfer the 2nd electron to another F Note that the coefficients CaF2 are the same as those in the balanced formula!

  16. 1.1 Forming and Representing Compounds • Bonds between ions, formed by electron transfer are called ionic bonds • Do question 6b, d, e on page 21

  17. 1.1 Forming and Representing Compounds • Molecular compounds – sharing electrons • Bonds formed by sharing electrons called covalent bonds • Covalent bonds exist between non-metallic atoms in a molecule and between atoms within a polyatomic ion (e.g. SO42-)

  18. 1.1 Forming and Representing Compounds • Drawing Lewis Structures for Molecular Compounds: • Start by counting the number of bonding electrons on each atom involved • The atom with the most bonding electrons is the central atom • Sometimes there will be more than 1 central atom • Start with the central atom, bonding other atoms to it, until it has the stable octet

  19. H 1.1 Forming and Representing Compounds • Examples: • H2O • • • • O H• • • 1 bonding electron 2 bonding electrons:central atom • Note that O gets the stable octet H ends up with 2 – that’s what it needs • H• • • O • •

  20. H 1.1 Forming and Representing Compounds • It is unnecessary to put circles around the atoms as your text does – I find they just make the diagrams messy and hard to read – even when typed out • • • • O   • H• •

  21. 1.1 Forming and Representing Compounds • A correct Lewis structure will always have a stable octet for every element other than H • When doing these diagrams, make electrons on adjacent atoms look different – different colour, different symbol, etc. Even if the atoms are atoms of the same element!

  22. H H • • • • • H C • H C • • • • • • • • H H 1.1 Forming and Representing Compounds • Further example: • C2H6 • H• C • • • central atom Note that both C’s get the stable octet Also note that I have made electrons on the 2 C’s different colours, since they are adjacent, and all H’s the same since they aren’t adjacent

  23. 1.1 Forming and Representing Compounds • Do BLM 1.1.4B – Lewis Diagrams of Molecules up to and including question 10 • Try question 11, but you will likely have some difficulties – I will show how to do this type with next set of examples

  24. stable octet on everything! •• Cl •• •• • • •• •• • • • Si • F • C • F • • • • • • •• • • •• Cl •• •• •• 1.1 Forming and Representing Compounds • Extra examples: • SiCF2Cl2 – always start with most symmetric arrangement • C • • • •• • Cl • Si •• •• • • • • • •• •• •• • • Si • • F F • C • F • • • • • • • •• • •• • • •• •• • Cl •• • • Cl • • •• •• Si and C don’t have stable octet. Share the other 2 electrons (circled in grey)

  25. •• • •• • P • Si • • H • P • • • • • 1.1 Forming and Representing Compounds • Example: HSiP • •• • • Si H • P • • • • • • Si • Si has 6; P has 6 - share more electrons! • H• Single, double, and triple covalent bonds are possible Experimental evidence indicates they actually exist

  26. •• O • • • • •• •• •• S O O • • • • • • • •• •• • • O • • • •• 1.1 Forming and Representing Compounds • Example: SO42- (polyatomic ion) – you will not be required to do these but I want to show you an illustration • extra electron • • Note: you could draw a Lewis Diagram with the extra electrons in other spots, but they there has to be 2 extra electrons to make SO42- stable

  27. 1.1 Forming and Representing Compounds • Extra Practice – Practice Problems 7 – 17, page 27 • Answers – page 728 – They don’t use colours or symbols to differentiate electrons, but this is a requirement for you

  28. •• •• •• •• •• •• •• •• O S O  •• 1.1 Forming and Representing Compounds • Resonance – page 27 actual structure is a cross between the two •• •• •• •• •• •• •• •• O S O •• Experimentally, bonds are identical These diagrams show resonance more clearly than the ones in your text You will not be required to draw resonance diagrams, but you should know the concept and be prepared to answer multiple choice questions concerning it

  29. + H •• • • H N H • • • • H 1.1 Forming and Representing Compounds • Coordinate covalent bonds – bonds where both electrons come from the same atom • e.g. NH4+

  30. 1.1 Forming and Representing Compounds • Structural formula – not the actual structure • Based on Lewis diagram – electron pairs are dropped and bonds replaced by lines H H H H • • • • • H C • H C • • H C H C • • • • • • H H H H

  31. 1.1 Forming and Representing Compounds •• Cl •• •• Cl • • •• •• • • • Si • F F • C Si C • F F • • • • • • •• • • •• Cl •• Cl •• •• •• • Si • • Si H • P H P • • • •

  32. 1.1 Forming and Representing Compounds • Structural formulas are very useful and we will use them, but don’t confuse them with molecular structure – Chapter 2 • Do question 8a, c, e, f, g, h, j page 31 – Lewis Diagram and structural formula for each • Modeling – chapter 2, after VSEPR theory

  33. 1.1 Forming and Representing Compounds • Metallic Bonding – theory now, relationship to properties in Chapter 2 • Metallic Bonding Handout - review Fig 1.16, page 32 shows Mg atoms having released their valence electrons to the electron cloud to become Mg ions in the metal crystal

  34. 1.1 Forming and Representing Compounds • Remember, metals do not exist as individual atoms – if they did they would be gasesThey exist as metal ions in a crystal of delocalized valence electrons • Analogy: Rice Krispee squares

  35. 1.1 Forming and Representing Compounds • Nomenclature Review: Question 9, page 34 • Identify each as ionic or molecular, and provide the correct name

  36. 1.2 The Nature of Chemical Bonds • Electronegativity: relative measure of an atom’s ability to attract shared electrons in a covalent bond Generalizations? metals low; non-metals high F highest; Fr lowest  from left to right  from top to bottom noble gases have none – not zero

  37. 1.2 The Nature of Chemical Bonds • Note 2 things about this diagram: atoms get smaller as you go across a row from left to right!! (but …….) electronegativity is inversely related to atomic size – small atoms  large electronegativity; large atoms  small electronegativity (Noble Gases excepted) fig 1.19, page 37

  38. 1.2 The Nature of Chemical Bonds • Atomic radius  as you move left to right because valence electrons for a given period all occupy same set of orbitals – outer radius is ~ fixed – but, as you move left to right nuclear charge  attracting valence electrons closer (attractive force and ) • Because of the small size and large charge, F has the highest electronegativity (attraction for shared electrons • Atomic radius  as you move from top to bottom since each period is filling of a new set of orbitals with larger radius

  39. 1.2 The Nature of Chemical Bonds page 38

  40. 1.2 The Nature of Chemical Bonds • Atoms can share electrons or transfer electrons to obtain the stable octet electron configuration • If electronegativities are close or equal you can predict electrons will be shared • If electronegativities differ by more than a certain limit you can predict electrons will be transferred • But, ………..

  41. 1.2 The Nature of Chemical Bonds • non-polar covalent bond: a covalent bond where electrons are shared equally • polar covalent bond: a covalent bond where electrons are shared unequally • a polar covalent bond is said to have a bond dipole – since electrons are shared unequally, one end of the bond is partially + (δ+); other end partially – (δ-) • there are 2 poles ……

  42. 1.2 The Nature of Chemical Bonds • Bond dipoles also designated by symbolExamples 10a, 11a, page 41 • Examples 10d, 11d, page 41 δ+ δ- C - F ΔEN = 4.0 – 2.6 = 1.4 polar covalent* ΔEN = 3.4 – 1.9 = 1.5polar covalent according to chart* we would normally call this ionic (continuum) Cu - O *chart, page 40 (next slide)

  43. know this number! 1.2 The Nature of Chemical Bonds Page 40 do questions 10 and 11 b, c, e, f, h, page 41

  44. 1.2 The Nature of Chemical Bonds • Another picture – page 88

  45. 1.2 The Nature of Chemical Bonds • Bonding Song

  46. 1.2 The Nature of Chemical Bonds • Review: Chapter 1 Review, page 44 good questions:1, 3 (what does each component of the diagram represent?), 7*, 8, 12, 14, 15 a, b, d, i, j, 16 (all but a and h), 17, 18, 19-24 (for question 24 draw bond dipole arrows for polar covalent bonds) * no multiple bonds here – also do C2HF, N2

  47. 1.2 The Nature of Chemical Bonds