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: NO 2 PCl 3 CaI 2

2. Chapter 8

3.  Covalent compounds consist of what?  Only nonmetals  When naming, we use …  Prefixes: mono, di, tri, tetra…  Prefix = number of atoms (subscript) N2O7N2O7  SF 6

4.  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.  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.  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.  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  CH 4  CCl 4

9.  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.  PH 3  H 2 S  SiH 4  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: O 2 N 2 F 2  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: C 2 H 6 C 2 H 4 C 2 H 2  Keep in mind how many times each element wants to bond

15.  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.  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.  PO 4 3- 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.  ClO 4 -  NH 4 +  CO 3 2-  H 3 O +  sulfite

19.  H 2 SO 4  CH 3 OH  HCN

20. Warm up: Name and draw the Lewis structures for the following compounds H 3 P CS 2 N 2 H 2

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

22.  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.  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 NO 3 -

24.  How many resonance structures do each of these have?  O 3  NO 2 -  SO 2  CCl 2 O

25.  Sometimes an atom may not obey the octet rule  Odd number of valence electrons (NO 2 )  Fulfill the octet of the “outer” atoms  Less than 8 electrons present around an atom (BH 3 )  Compounds with Be or B  Tend to be very reactive  Coordinate covalent bond – when one atom donates both electrons in a shared pair (BH 3 + NH 3 )

26.  Draw the Lewis structure for SO 3 and draw its resonance structures  Draw the Lewis structure for ClF 3

27.  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  PCl 5  SF 6  Let’s look at H 2 SO 4 again  The S-O bonds have been experimentally determined shorter than single bonds

28.  ClF 5  More than an octet on chlorine  ICl 4 -1  More than an octet on iodine  BeH 2  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.  Draw the Lewis structures for ammonia (NH 3 ) and the ammonium ion

30.  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 NH 3 and NH 4 + 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.  Draw the structures and determine the FC for each atom Cl 2 O SO 2 AsF 3

33.  V alence S hell E lectron P air R epulsion – 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.  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. Linear Bent Trigonal planar Tetrahedral Trigonal pyramidal Trigonal bypramidal Octahedral 107.3 o 104.5 o 120 o 109.5 o 90 o / 120 o 180 o 90 o

36.  If the bond is not lying in the plane, then you use either dashes or wedges

37.  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.  NCl 3  OCl 2  HOF  NHF 2  CO 2  H 2 Se  CH 2 O  NH 4 +1

39.  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 10 th

40.  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? sp 3

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

43. Regions of high e- density VSEPR shapeHybridization 2Linearsp 3Trigonal planarsp 2 4Tetrahedralsp 3 5Trigonal bipyramidalsp 3 d 6octahedralsp 3 d 2 When giving the hybridization, you are generally talking about the hybridization for the central atom

44.  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 sp 3

45.  NCl 3  OCl 2  HOF  NHF 2  CO 2  H 2 Se  CH 2 O  NH 4 +1

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

47.  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.  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 DifferenceBond Type Less than 0.4Nonpolar covalent 0.5 to 1.9Polar covalent Greater than 2.0Ionic  What kind of bond would carbon and oxygen form?  Phosphorus and fluorine?  Chlorine and chlorine?

50.  Draw the Lewis structure, determine the shape and hybridization for the following: BF 3 SF 4 PF 6 -

51.  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.

52.  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?  You have to look at the polarity of each bond, and look at the overall molecule to determine if it is polar

53.  Determine if the following molecules/ion are polar: NCl 3 H 2 S CS 2 SF 6 If the bonds are polar, it could be polar or nonpolar, check the structure

54.  Solubility (what is this?) is determined by polarity  What is the universal solvent?  Are most substances polar or nonpolar?

56. Determine the more polar molecule in each pair: methyl chloride (CH 3 Cl) or methyl bromide (CH 3 Br) water or hydrogen sulfide (H 2 S) hydrochloric acid or hydroiodic acid boron trihydrideORammonia silicon tetrabromide ORHCN

57.  What were the properties of ionic compounds in terms of conductivity, melting point and solubility?  High melting point, conducts (when dissociated), and soluble in water (meaning ionic compounds are what?)  What are properties of covalent?  Many covalent compounds exist as liquid or gas  Which type is more strongly held together?

58.  What are intermolecular forces? (interstate)  Forces that hold one molecule to another 3 Types: Hydrogen bonding Dipole-dipole Dispersion/London forces  In the solid/liquid state (not concerned with gaseous state – why?)

59.  Dipole – contains oppositely charged regions (partial charges)  Results from the attraction between the partial positive end of one molecule and partial negative end of another molecule

60.  Also known as induced dipole forces  animation animation  Occur between nonpolar molecules with no permanent dipoles  Result from a temporary shift of electrons, and dipoles are instantaneously created  Ex. 2 chlorine molecules

61.  Occur between hydrogen and O, N or F  Due to their high electronegativities it makes H more partially positive  Causes these compounds to have higher boiling points

62. What is the strongest intermolecular force present for each of the following compounds? 1) water 2)carbon tetrachloride 3)ammonia 4)carbon dioxide 5)phosphorus trichloride 6)nitrogen 7)ethane (C 2 H 6 ) 8)acetone (CH 2 O) 9)methanol (CH 3 OH) 10)borane (BH 3 )

63. 1) water hydrogen bonding 2)carbon tetrachloride London dispersion forces 3)ammonia hydrogen bonding 4)carbon dioxide London dispersion forces 5)phosphorus trichloride dipole-dipole forces 6)nitrogen London dispersion forces 7)ethane (C 2 H 6 ) London dispersion forces 8)acetone (CH 2 O) dipole-dipole forces 9)methanol (CH 3 OH) hydrogen bonding 10)borane (BH 3 ) dipole-dipole forces

64.  Grab a chemistry book, and work on the following questions – p. 27483, 85, 89, 96, 98, 101, 108, 112, 114, 120, 127 Be sure to look through my powerpoints and study guide on my website

65.  Name the following compounds:  ZnCl 2  KNO 3  H 2 S  NF 3

66.  Name and draw the Lewis structures for the following compounds:  CS 2  PH 3  CCl 4

67.  Write the formulas for the following compounds:  Aluminum sulfate  Iron (III) phosphide  Hydronitric acid  Nitrous acid  Dicarbon trisulfide  Go ahead and take out the worksheet with the PT with electronegativities from yesterday

68.  Name the following acids: H 3 N H 3 SO 3 H 2 Se