Tests are not graded yet Turn in your project up front and work on warm up:

1. Tests are not graded yet Turn in your project up front and work on warm up: • Write the molecular formula for: Trinitrogen hexoxide Aluminum nitride Copper (II) sulfate Write the names for: NO2 PCl3 CaI2

2. Covalent bonding Chapter 8

3. Covalent compounds • Covalent compounds consist of what? • Only nonmetals • When naming, we use … • Prefixes: mono, di, tri, tetra… • Prefix = number of atoms (subscript) • N2O7 • SF6

4. 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 exception to this rule? • Which group do all elements want to be like?

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

6. 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

7. 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

8. 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 will equal how many times an element will bond • CH4 • CCl4

9. 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”

10. 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

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

12. 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)

13. Draw: O2 N2 F2 • 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?

14. Warm up: Draw the Lewis structures for the following: C2H6 C2H4 C2H2 • Keep in mind how many times each element wants to bond

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

16. Bonding and energy • In chemical reactions, bonds are broken and formed • Breaking bonds _____________ energy • Requires (breaking a stick) • Forming bonds _____________ energy • Gives off (Aladdin) • If more energy goes in, then it is _______________ • Endothermic • If more energy is given off, then it is ___________ • Exothermic

17. 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

18. ClO4- • NH4+ • CO32- • H3O+ • sulfite

19. Try these bad boys… • H2SO4 • CH3OH • HCN

20. Warm up: Name and draw the Lewis structures for the following compounds H3P CS2 N2H2

21. HONC • H – 1 time • O – 2 times • N – 3 times • C – 4 times • Lowest electronegativity element goes in the center

22. Diatomic molecules • Look at the word… • Molecules that contain how many atoms? • My fish’s name will help you know these • 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.

23. Resonance • What does it mean when something resonates? • To vibrate or sound, especially in response to another vibration • Resonance structures are different ways to draw Lewis structures for a molecule or ion • Only the arrangement of the electrons is changed • Let’s draw the structure for NO3-

24. How many resonance structures do each of these have? • O3 • NO2- • SO2 • CCl2O

25. Exceptions to the octet rule • Sometimes an atom may not obey the octet rule • Odd number of valence electrons (NO2) • Fulfill the octet of the “outer” atoms • Less than 8 electrons present around an atom (BH3) • Compounds with Be or B • Tend to be very reactive • Coordinate covalent bond – when one atom donates both electrons in a shared pair (BH3 + NH3)

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

27. Exceptions to the octet rule • Expanded octet: happens with elements in period 3 and below – d orbital electrons can hold more than 8 • Generally, the central atom gets the extra electrons • PCl5 • SF6 • Let’s look at H2SO4 again • The S-O bonds have been experimentally determined shorter than single bonds

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 structures for ammonia (NH3) and the ammonium ion

30. Formal charge • The hypothetical charge on an atom in a covalently bonded molecule • Helps to determine the best Lewis structure • Want to keep the formal charge low – most stable structure FC = (# valence e-) – [(# of bonds) + (# of unshared e-)] In a molecule, the sum of the formal charges (for every atom in the molecule) is zero In a polyatomic ion, the sum is equal to the charge

31. Use the structures for NH3 and NH4+ from the warm up • Determine the FC for each nitrogen and hydrogen in both structures • Write the value next to the atom; if there is no number, it is understood to be zero • Draw the structure for NOCl • There are 2 possibilities, one is more preferred • Draw the structure for sulfate

32. Practice • Draw the structures and determine the FC for each atom Cl2O SO2 AsF3

33. 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

34. Detemining shapes • On a separate sheet, draw the Lewis structures for each of the compounds on the handout • Let’s see how many bonded pairs there are, and how many lone pairs on the central atom there are • Don’t fill in the picture column or angle column yet

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

36. Drawing Lewis structures – 3D • If the bond is not lying in the plane, then you use either dashes or wedges H C H H H

37. Why is water bent? • When electrons are bonded, think of them as “trapped” between the 2 atoms, therefore occupying less space • Lone pairs occupy more space, therefore causing the bonded electrons to repel (and bend the molecule)

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

39. Extra Credit • Pick one of the VSEPR shapes and build a molecule • Include: label the type, an example of a specific molecule (none that are on the table), 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 next Wedn. Feb 10th

40. Hybridization • Hybrid – when 2 things combine and have properties of both • When atoms bond, they want to arrange their orbitals to have lowest energy possible • Hybridization – describes the arrangement of the orbitals • Hybrid orbitals – combined orbitals; intermediates between orbitals • between s and p lies the hybrid orbital sp

41. Draw the orbital diagram for Carbon • From this, it looks as if there are only 2 places for electrons from another atom to pair up (in the p orbital), but how many times does carbon like to bond? sp3

42. Warm up Write the formulas for the following compounds: • Aluminum sulfate • Iron (III) phosphide • Hydronitric acid • Nitrous acid • Dicarbon trisulfide

43. Cl Be Cl • When giving the hybridization, you are generally talking about the hybridization for the central atom

44. Hybridization • Generally, the # of things you are bonded to = the number of hybrid orbitals • Bonded to 2 things = sp • Lone pairs(on the central atom) occupy hybrid orbitals as well • Ex: draw the Lewis structure for water • Those 2 lone pairs count towards the hybrid orbitals, so water is sp3

45. Determine the hybridization… • NCl3 • OCl2 • HOF • NHF2 • CO2 • H2Se • CH2O • NH4+1

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

47. Polarity • Influenced by the electronegativities of atoms in a molecule • What is electronegativity? • An atom’s attraction for electrons when in a bond • What is the trend for electronegativity? (remember shielding and nuclear strength) • Who has the highest electronegativity value?

48. 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

49. 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?

50. Warm up • Draw the Lewis structure, determine the shape and hybridization for the following: BF3 SF4 PF6-