Covalent bond lewis structure Molecule electronegativity Intermolecular forces sigma bond

1. Covalent bond lewis structure Molecule electronegativity Intermolecular forces sigma bond Pi bond orbital Chemical bond exothermic reaction Endothermic reaction bond length Bond dissociation energy octet rule Resonance VSEPR model Polar covalent bond structural formula Vocab quiz moved to tomorrow - you have 20 minutes to work on them today

2. Warm up: • Write the molecular formula for: Trinitrogen pentoxide Aluminum nitride Copper (II) sulfate Write the names for: NO2 PCl3 CaI2

3. Covalent bonding Chapter 8

4. Covalent compounds • Covalent compounds consist of what? • Only nonmetals • When naming, we use … • Prefixes: mono, di, tri, tetra… • Prefix = number of atoms (subscript) • N2O7

5. Covalent compounds • Why are there no charges (like in ionic compounds)? • In ionic compounds, electrons are _______________, so atoms gain or lose charge • In covalent compounds, electrons are _____________, so no charges are formed • What does the octet rule state? • In order to be stable, an atom wants a full outer shell (which generally means 8 valence electrons) • Which nonmetal is the only exception to this rule? • Hydrogen – how many does he need to be full? transferred shared

6. What is a covalent bond? • When neither atom wants to give up their electrons, they will just share • When 2 electrons are shared between atoms, they form a single bond • When 2 or more atoms bond covalently, this is called a molecule

7. Why do they share? • Consider ionization energy and electronegativity– when 2 elements are near each other on the periodic table, these values will be very near each other • Ionization energy • Energy required to remove an electron • Electronegativity • How well an element attracts electrons in a bond

8. If both atoms have very similar strengths (for holding on to their electrons) then…. • Neither one will be strong enough to take electrons away from the other

9. Drawing covalent compounds • Lewis structures – using electron dot diagrams, shows the arrangement of the atoms in a molecule • How many valence electrons does carbon have? How many more electrons does it need to be “happy”? • How many times do you think carbon will bond? • How about hydrogen? Oxygen? • Generally, the # of “missing” electrons is equal to the number of times an element will bond • HONC 1234

10. Drawing Lewis structures • Calculate the number of valence electrons • Arrange the atoms in the molecule • Generally, the atom you have one of will go in the middle • Hydrogen only bonds once, bonds on the outside • How many times will carbon bond? Oxygen? (look at their valence electrons) • Put pairs of electrons between the central atom and all of the outer atoms • Put electrons to fill the central atom • Put remaining electrons around outer atoms • Check to see that every atom is “happy”

11. Take a minute and look over your vocab… • When you are finished with the quiz, bring it up front and take out your notes from yesterday

12. CCl4 • PH3 • H2S • SiH4 • When 2 electrons are shared between atoms, you draw a line to show the bond • All other electrons that are not shared are called lone pairs and are included in the structure

13. Single covalent bonds are also called sigma bonds • Orbitals – the area where you will most likely find an electron • When these orbitals overlap, they form a sigma bond (σ)

14. Let’s try carbon dioxide… • Sometimes, atoms may share more than 2 electrons • If 4 electrons are shared, how many bonds would there be? • This is called a double bond • How many electrons would a triple bond share? • Double or triple bonds consist of sigma and pi bonds (π)

15. A word of caution!! • Don’t start using double bonds in every molecule, it is a last resort • Bond everything first, use up the rest of your electrons, THEN see if you need to double or triple bond • Try SO2

16. Warm up • Draw the Lewis structures for SO2 and OF2

17. Diatomic molecules • Look at the word… • Molecules that contain how many atoms? • H. BrONClIF • In nature, when these elements are not bonded to another element, they like to exist with 2 of themselves. They are more stable that way.

18. Draw: F2 O2 N2 • What do you notice about the bonds? • Bond length : the distance between two bonding nuclei • Which of these 3 do you think would have the shortest bond length?

19. Bond length and energy • As the number of bonds increases, the bond length becomes shorter • Which bond would be the strongest? • Shorter = stronger • Bond dissociation energy : energy required to break a bond in a molecule • What is the relationship between bond length and dissociation energy? • Shorter = more energy

20. Lewis structures of polyatomic ions • PO43- what is this called? When an ion has a charge, that means it has lost or gained ______________ What has phosphate done? Start the lewis structure like we did for the others – add up all valence electrons Now we have 3 extra electrons

21. ClO4- • NH4+ • CO32- • H3O+ Be careful ! a negative charge means we are gaining electrons, while a positive charge means we have lost electrons!

22. Try these bad boys… • H2SO4 • C2H4 • N2H2

23. Warm up • Determine the chemical formula and draw the Lewis structures for the following compounds or ions: • Carbon disulfide • Phosphorus trihydride • Sulfate ion

24. Resonance • What does it mean when something resonates? • Resonance structures are different ways to draw Lewis structures for a molecule or ion • Only the arrangement of the electrons is changed – keep the skeleton structure • Let’s draw the structure for NO3-

25. How many resonance structures do each of these have? • O3 • NO2- • SO2 • CO3-2

26. Exceptions to the octet rule Sometimes an atom may not obey the octet rule 1. Odd number of valence electrons (NO2) 2. Less than 8 electrons present around an atom (BH3) 3. Expanded octet: happens with elements in period 3 and below – d orbital electrons • Generally, the central atom gets the extra electrons • PCl5 • SF6

27. Look over your notes for the quiz… • If you have a 75 or lower on your progress report, you must get it signed by a parent and returned to me by Wednesday • Try the following, and refer to notes from yesterday: ICl3

28. What is “wrong” with these structures? • ClF5 • More than an octet on chlorine • ICl4-1 • More than an octet on iodine • BeH2 • Less than an octet - Beryllium and boron generally follow the less than 8 exception • NO • Odd number of valence - Nitrogen generally takes the odd number of electrons

29. Warm up Draw the Lewis structure for the following: ICl3 BeCl2 NS2 Any progress reports?

30. VSEPR theory • Valence Shell Electron Pair Repulsion – used to determine the shape of a molecule • What determines how a molecule will arrange itself? • What part of the atom are we generally concerned about?... • ELECTRONS • Something to keep in mind: lone pair electrons occupy more space than bonded electrons

31. Building VSEPR Models • Draw the Lewis structures for each of the compounds on the handout (1st one is already done) • Let’s see how many things the central atom is bonded to, and how many lone pairs on the central atom there are

32. Bond Angles –between atoms in a molecule 180o 104.5o 120o 109.5o 107.3o 90o/ 120o 90o

33. Determine the shapes… • NCl3 • OCl2 • HOF • NHF2 • CO2 • H2Se • CH2O • NH4+1

34. Extra Credit • Pick one of the VSEPR shapes and build a molecule • Include: label the type, an example of a specific molecule, the angle between the atoms, represent lone pairs (if there are any) • Use anything you would like to build this – no drawings, and the model must be an accurate representation of the shape • Due Friday, Feb. 12th

35. Polarity • If something is polar, it means it has opposing ends • Need to know electronegativity and shapes

36. Polarity • Influenced by the electronegativities of atoms in a molecule • What is electronegativity? • How well an atom attracts electrons in a bond • What is the trend for electronegativity? (remember shielding and nuclear strength) • Increases up and over • Who has the highest electronegativity value? • Fluorine

37. What do these numbers tell you? • Ionic: Look at the electronegativities of Na and Cl – who has more attraction for the electrons? • Covalent: look at the values for the nonmetals • Polar covalent – unequal sharing of the electrons in a bond • Nonpolar covalent – equal sharing of electrons in a bond

38. Electronegativity Difference Bond Type Less than 0.4 Nonpolar covalent 0.5 to 1.9 Polar covalent Greater than 2.0 Ionic • What kind of bond would carbon and oxygen form? • Phosphorus and fluorine? • Chlorine and chlorine?

39. Arrange in order of increasing polar character C – O Si – O Ge – O C – Cl C - Br

40. Draw the Lewis structure for water • What is water’s shape? • Who is stronger? • Who will the electrons be closer to? • This makes partial charges.

41. Determining polarity • Draw carbon tetrachloride and label the partial charges • Compare carbon tetrachloride’s structure to water’s • Polar molecules are asymmetric, while nonpolar are symmetrical • Which one of these would you consider symmetrical?

42. Determining polarity • Determine if the following molecules/ion are polar: H3O+ NCl3 H2S oxygen CF4 CS2 SF6

43. LIKE DISSOLVES LIKE • Solubility (what is this?) is determined by polarity • What is the universal solvent? • Are most substances polar or nonpolar?

45. Practice – there will be compounds on test • Write the formulas for the following compounds: • Aluminum sulfate • Iron (III) phosphide • Hydronitric acid • Sulfurous acid • Dicarbon trisulfide

46. Review • Grab a chemistry book, and work on the following questions – p. 274 83, 85, 89, 96, 98, 101, 108, 112 (don’t worry about hybrid orbitals), 114, 120, 127 Be sure to look through my powerpoints and study guide on my website

47. Warm up: • Draw the Lewis structure for SO3 and draw its resonance structures • Draw the Lewis structure for ClF3

48. Warm up: • Name the following compounds: • ZnCl2 • KNO3 • H2S • NF3