CHOPPED-UP CHAPTER. not too chopped up, but enough

Transcription

1 CHOPPED-UP CHAPTER not too chopped up, but enough

2 chapter 12 chemical bonding

3 what do you know? 1.When representing a water molecule as H-O-H, what do the lines between the letters symbolize? 2.Why is the term sodium chloride molecule incorrect?

4 here we will... learn about ionic and covalent bonds and explain how they are formed

5 essentially everything we see around us - natural or man-made, cmpds or elements - has atoms bonded with atoms hard or soft? solid, liquid, or gas? these properties are determined by bonding molecular structures play the major role in chm rxns all around us, and and the way they re bonded determines structure

6 12.1 types of chemical bonds bond: a force that holds groups of two or more atoms together and makes them function as a unit they can do this in several ways! ready?

7 remember table salt? NaCl when the water in salt water evaporates away, the Na+ get together with the Cl- to form NaCl

8 this closely packed collection of oppositely charged ions is an example of ionic bonding

9 how strong are ionic bonds judging by this?

10 these substances are called ionic compounds, and they can be made like this

11 remember: metals like to lose e- s and nonmetals gain so there can be a transfer here of electrons and the result is two things that love each other :)

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14 M + X M + X - M = metal, X = nonmetal, M+ X - is the ionic cmpd (e.g. NaCl, MgO, KBr ) but what if the atoms are identical? how do they bond together?

15 when these two critters get close, the e- of one is attracted to the p of the other; the e - (or p) of one is repelled by the e - (or p) of the other

16 if they are most comfy at a distance where their e- shells overlap - ba da bing! - we gots us a bond! most comfy here

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18 mini-summary this is a covalent bond because outer shell (valence) electrons are shared by both atoms

19 bigger-than-mini summary

20 ionic bonding (transferring e - s) and covalent bonding (sharing e - s) are the extremes b/t these 2 extremes are

21 polar covalent bonds! here one atom is not strong enough to strip the electron off the other, but it can hog it the hog gets a - to show it is almost (not quite) a negative ion; the loser gets a + but how much do these critters like electrons???

22 quick quiz ionic, covalent, both, neither 1.This one involves the transfer of electrons. 2.Valence electrons are involved here. 3.This almost always involves metals with nonmetals. 4.This one happens because of the overlapping of outer shells. 5.Here there is a force that holds groups of two or more atoms together and makes them function as a unit.

23 12.2 electronegativity when different nonmetals form a cov bond, they almost never share the electrons equally electronegativity = the relative ability of an atom in a molecule to attract electrons to itself

24 here are electronegativity values; see a general pattern?

25 F rules the en table! and the numbers generally decrease as you move away from him

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28 electronegativity ionization energy size

29 bottom line: even though they share electrons they aren t very nice about the sharing electronegativity will tell us who the selfish ones are

30 the polarity of a bond depends on the two battling for the electrons in fact, the difference in electronegativity is the Big Deal here: great differences in en = more ionic sort of bonds small differences = polar covalent bonds no difference = (nonpolar) covalent

31 en bond type 0 - ~0.4 nonpolar covalent ~0.4 - ~2 polar covalent > ~2 ionic e.g. here is a polar covalent situation

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33 but don t think it s so black and white; it s more like this

34 frinking frap time Write down any patterns you see here. Take your time!

35 now we try some examples

36 quick quiz 1 Classify these as I, PC, or NPC: PG KF 2. O 2 3. IC 4. HBr 5. NaF 6. N 2 Δ = 3.2 Δ = 0 Δ = 0.5 Δ = 0.7 Δ = 3.1 Δ = 0 I NPC PC PC I NPC

37 quick quiz 2 Organize these from least polar to most 1. H-H Δ = O-H 3. C -H 4. S-H Δ = 1.4 Δ = 0.9 Δ = PG F-H Δ = 1.9 5

38 2, 3, 8, 9, 11 12EOCs

39 12.3 bond polarity and dipole moments if the molecule actually ends up having a partial positive side and a partial negative side we say it has a dipole moment an arrow tells us where the negative side is

40 it always happens in diatomic molecules with unequal sharing; it can also happen in polyatomic molecules like water:

41 this unequal distribution of charge in a water molecule is unbelievably important to life

42 it can surround and dissolve ions (more on that later)

43 they can attach to each other and stick (more later) which is why it is a liquid and a solid on earth

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45 here two polar bonds cancel each other out to give a nonpolar molecule (red O s pull away electrons from the weaker C)

46 do you see that this molecule s polar bonds do not get cancelled out, so the molecule overall is polar???

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48 General Rule If the molecule is asymmetric, there is probably some uneven distribution of electrons and it will be polar. If the molecule is symmetric, probably any polar bonds cancel each other out and the whole thing is nonpolar.

49 what if...? We use differences in electronegativity to account for certain properties of bonds. What if all atoms had the same electronegativity values? How would bonding between atoms be affected? What are some differences we would notice?

50 quick quiz 1. How are ionic bonds and covalent bonds different? 2. How does a polar covalent bond differ from a nonpolar covalent bond? 3. Match the compounds HF, NaCl, and O2 with the following figures: 4. For each of these draw an arrow in the direction of its dipole moment. a. H-Cl b. B-H c. H-I d. Br-Br e. C-O PG 362

51 walk away with... Chemical bonds hold groups of atoms together to form molecules and ionic solids Bonds are classified as: ionic: formed when one or more electrons are transferred to form positive and negative ions covalent: electrons are shared between atoms electronegativity is the relative ability of an atom to attract the electrons shared with another atom in a bond

52 1, 2, 3, 4, 6, 8 12EOCs

53 here we will... learn about stable electron configurations learn to predict formulas of ionic compounds learn about the structures of these compounds understand factors governing ionic size

54 12.4 stable electron configurations and charges on ions what you shd get out of this: these critters will gain or lose electrons until they look like a noble gas...

55 after this and lots more observations we can say this: in almost all stable compounds of the representative elements, all of the atoms have achieved a noble gas configuration there is just something special about getting a full outer shell!

56 so when Ms react with NMs they ll exchange e - s until they look noble [important note: they don t become noble gases, they just have an electron config that looks like one] when NMs react with NMs they ll share until they look noble

57 this is how they all look when they achieve nobility; look familiar?

58 we can summarize observations like this with a modified PT: nonmetals NOBLE metals

59 predicting formulas of ionic compounds now we can have a CHM A review with more depth in the old days we saw that Ca ion is +2, and O has a -2, to form CaO now we know Ca with [Ar]4s 2 will lose two electrons (to look noble) thus Ca2+, and

60 O is [He]2s 2 2p 4 so it needs how many to look noble? 2! it can take the two electrons from Ca, and now be O 2- so both get a noble appearance and the result is Ca 2+ and O 2- to form CaO got it!?

61 so what happens when Al metal is reacted with oxygen? Al has its 3 valence electrons, and O only needs 2 so two Al atoms can provide enough electrons for three O s (6 total) so, Al 2 O 3

62 12.5 ionic bonding and structures of ionic compounds when we learned that sodium chloride is NaCl we really were talking empirical (simplest) formula in reality there are b-zillions of Na s packed together with bzillions of Cl s in a 1:1 ratio

63 here is the structure for LiF again, the LiF just tells us the ratio of these two is 1:1 they are closely packed and tough to move around (making it brittle)

64 it s easier to see the structure s geometric shape with the wire frame model

65 you may have noticed that the positive ions are always dinkier than the negative ion; why? remember the M loses its whole outer shell to form an ion; and the NM gains electrons to look noble (getting bigger as a result)

66 see an obvious pattern?

67 what if...? Ions have different sizes than their parent atoms. What if ions stayed the same sizes as their parent atoms? How would it change the structure of ionic compounds? (hint: look how nice and closely packed they are now.)

68 ionic compounds containing polyatomic ions [this always causes a bit of confusion so pay attention here] the polyatomic ions (like NH 4 + or NO 3 -) are true ions with real charges and can get together with other regular ions, but! they are covalently bonded within themselves

69 the polyatomic ion acts like a single welded-together ion so when ammonium nitrate (NH 4 NO 3 ) dissolves in water it splits into NH 4 + and NO 3 - only; each of the ions stays intact, like a charged family OK?

70 quick quiz 1. Why do metals lose electrons to form ions, and when do they stop losing them? 2. Why does oxygen form O 2- and not O 3-? 3. Predict the formula for each combo. a. Mg, S b. K, Cl c. Cs, F d. Ba, Br

71 walk away with... In stable compounds, atoms tend to achieve the electron configuration of the nearest noble gas atom In ionic compounds: nonmetals tends to gain electrons to reach noble gas configuration metals tend to lose electrons to get there Ions group together to form cmpds which are electrically neutral

72 12EOCs 20 (only second part of question) 23, 24, 28

73 here we will... learn how to write lewis structures learn to write lewis structures for molecules with multiple bonds

74 12.6 lewis structures remember: bonding involves the valence electrons of atoms!!! the valence are transferred in ionic bonding and shared in covalent the Lewis structure is just a simple representation to show how the valence e - s are distributed around a molecule

75 it was developed 100 years ago by G.N. Lewis chemists had figured out by then that the valence electrons are used to make everyone look noble he used dots to show valence e - s, sorta like

76 this!

77 this! see a pattern???

78 making ionic cmpds is a breeze Ms lose all their valence electrons; NMs gain until they look noble

79 here s a visual way of looking at the reaction of Al with Br 2 in a Lewis way the real rxn is more complicated; this focuses on where the e - s are going

80 here s a visual way of looking at the reaction of Al with Br 2 in a Lewis way the real rxn is more complicated; this focuses on where the e - s are going

81 we write them like this:

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84 the covalent critters are a little more involved they still have to look noble and they share until all are happy let s start easy

85 H s with their lowly single valence only need one more to be dressed like helium, so they agree to share: see how we represent it with the dots? notice each H sees two electrons

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87 its desire for 2 total e - s (to look like He) is called the duet rule so whenever H is involved make sure it ends up with 2 electron dots only!

88 the second-row NMs want chemists call an s 2 p 6 look; i.e. they want a total of eight! called the octet rule

89 they will share two per bond like this: notice they both see 8 e - s

90 a shared pair is called a bonding pair not shared? = lone pair or unshared pair noble gases are not invited to these soirees ready for LD structures for bigger molecules? (answer: yes) first, the rules

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92 example draw the lewis structure for water 1. collect all the valence electrons H + H + O = = 8 2. give each bond 2 H:O:H 3. give the leftovers to the ones that need them.. H:O:H or with lines for bonds H O H......

93 what s all this then? why is the LD structure all crooked??? cool your jets! - all the requirements are fulfilled! H sees 2, O sees 8 there are several correct versions of most lewis structures; remember, it s just telling us where the electrons are not how the molecule is structured

94 example write the lewis structures for NH 3 and CH 4

95 quick quiz write the lewis structures for HCl PH3 H2S.. H Cl:.... H P H H.. H S H..

96 12.7 lewis structures of molecules with multiple bonds what happens if you run out of electrons before everything has its 8? e.g. CO 2 : C + O + O = = 16 start with O C O and you have 12 left but each O will need 6 and then you ve run out! what about poor carbon??? :(

97 how can you get carbon its 8? the O s agree to share their lone pairs and we make the single bond into a double bond! [think symmetrically!] you can form this:

98 you may have written this: this is a triple bond! the sharing of 3 pairs notice that any of these is correct according to the rules, but experiment shows the symmetric one to be the best when >1 structure can be drawn for a cmpd, that is resonance

99 what about lewis structures for ions? no problem! if it is a neg ion just add electrons to cover the charge if positive, subtract from total to cover the charge, like

100 write the lewis structure for CN- C + N +1 = = 10 C N is the base add electrons so everyone is happy but it must be written [:C:::N:] - this would not have worked w/o the extra electron

101 do lewis figures for: quick quiz 1 a) HF b) N 2 c) NH 3

102 quick quiz 2 do lewis figures for: d) CH 4 e) CF 4 f) NO +

103 some exceptions to the octet rule some atoms are too small to get a full eight; some are so big they can accommodate more e.g. little boron is happy with 6 or 8 Be can make cmpd with only 4 e - s DON T BE ASCARED!

104 quick quiz For each molecule below: a. ClF b. Br2 c. H2O d. O2 Give the sum of the valence electrons. Draw the Lewis Structure. Circle the octet or duet for each atom in the molecule.

105 walk away with...1 In covalent compounds, nonmetals share electrons so that both atoms achieve noble gas configs Lewis structures represent the valence electron arrangements of the atoms in a cmpd

106 walk away with...2 The rules for drawing Lewis structures recognize the importance of the noble gas electron configurations duet rule for hydrogen octet rule for most other atoms Some molecules violate the octet rule for the component atoms examples: BF3, BeCl2

107 N 2 NCl 3 BSF MORE! :)

108 CF 4 N 2F4 PBr 3 AND MORE!

109 NO 2 N 2H2 C 2H4 AND MORE!!!

110 33, 34, 35, 37 12EOCs

111 here we will... understand molecular structure and bond angles learn to predict molecular geometry from the number of electron pairs learn to apply the VSEPR model to molecules

112 12.8 molecular structure so far: where the electrons go now: how they look in 3D called molecular structure (or geometric structure) first the basic shapes then figuring out which of those shapes molecules take

113 this molecule has a bent shape with bond angles of ~105

114 some molecules have a bond angle of called linear structure

115 some molecules have a bond angle of called linear structure

116 there s a third shape called trigonal planar with 120 bond angles

117 there s a third shape called trigonal planar with 120 bond angles

118 one of the most important, though, is the tetrahedral this critter has four things sticking out of it in four symmetric directions the magic bond angle is but what determines what the shape will be?

119 one of the most important, though, is the tetrahedral this critter has four things sticking out of it in four symmetric directions the magic bond angle is but what determines what the shape will be?

120 39, 41 12EOCs

121 12.9 molecular structure: the vsepr model you can t believe the importance molecular structure has in life! one tiny change on a giant biomolecule can render it useless here we learn how to predict the basic structures of molecules based on those crazy electron pairs (shared and unshared)

122 we ll use a simple but effective intro to molecular 3D structure called VSEPR = valence shell electron pair repulsion model the name says it all; we re talking about shapes of molecules due to: 1. pairs of e - s (shared & lone) in the valence shell and 2. their dislike of each other

123 we want the pairs of e- s to get as far away from each other as they can we ll start with an easy one, and an exception to the octet rule, BeCl 2 the Lewis structure is like this: concentrate on the central atom, Be those shared pairs don t like each other and will move as far away as they can

124 the farthest two pair can get is 180 this forces BeCl 2 into a linear molecule

125 now BF 3 (another exception to the octet rule) has a central atom surrounded by three shared pairs how far away can they get? 120! [note: Lewis structure nothing like geometry of real thing; was never meant to be]

126 this molecule, shaped like a triangle and flat (aka planar) is called trigonal planar what about 4 pairs of electrons?

127 how do we build CH 4, a central atom with 4 electron pairs around it? this is real new and counterintuitive

128 most people when seeing the Lewis structure think it must be a cross but the electron pairs can get farther than 90! they can actually get away

129 a central atom with four electron pairs forms something called a tetrahedron

130 what if...? You have seen that molecules with four electron pairs, such as CH4, take on a tetrahedral shape. What if a molecule had six electron pairs like SF6? Predict the geometry and bond angles of SF6.

131 important distinction alert! we separate all the electron pairs, shared or alone, using vsepr, but! we are most concerned about where the atoms are as a result we will name the molecular structure based on where the atoms end up watch

132 first the rules 1. draw the Lewis structure! 2. count the electron pairs (bonded and unshared) and arrange them so they re as far away from each other as they can get 3. determine where the atoms go 4. name the thing from where the atoms are! now, some examples

133 predict the structure for NH 3 1. get the Lewis structure 2. see that there are four electron pairs around N; they would be placed like this

134 put the atoms in:

135 name that bad boy, remembering when you name it just worry about where the atoms are pairs are tetrahedral; but molecule is trigonal pyramid [not trigonal planar; the lone pair has pushed the H s down into a pyramid]

136 now water 1. get the Lewis structure 2. see that there are four electron pairs around O; they would be placed like this

137 put the atoms in:

138 name that bad boy, remembering when you name it just worry about where the atoms are pairs are tetrahedral; but molecule is bent (aka V-shaped) [not linear! the lone pairs pushed the H s over]

139

140 examples (molecular) Cl 2O molecule is bent NF 3 molecule is trigonal pyramid

141 examples (molecular) Cl 2O molecule is bent NF 3 molecule is trigonal pyramid

142 examples (molecular) H 2Se molecule is bent (V-shaped) ClO 4 - molecule is tetrahedral

143 examples (molecular) H 2Se molecule is bent (V-shaped) ClO 4 - molecule is tetrahedral

144 examples (just molecules) NH 4 + H 2 S BeF 2

145 examples (just molecules) NH 4 + molecule is tetrahedral H 2 S molecule is bent BeF 2 molecule is linear

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147 47, 48, 56 12EOCs

148 12.10 molecular structure: molecules with double bonds vsepr works with double bonds, too! we see in experiments that CO 2 takes on a linear look it s as if the double bonds of CO 2 (as far as being repulsive)! act like single bonds

149 if it helps look at single and double bonds as electron clouds of repulsion single! double! the double pair are restricted to an area just like a single pair (they re just a wee fatter)

150 it happens here, too, with ozone :O O=O: experiment shows that ozone is shaped like this: do you see why it is bent and not linear?

151 look at the center oxygen :O O=O: it has a lone pair, a single bond, and a double bond! this behaves as if there were just three areas of electrons thus,

152 conclusion: when using the vsepr to predict the geometry of a molecule treat the double or triple bond as if it were a single

153 quick quiz 1. How does a molecule s Lewis structure help determine its shape? 2. When is the molecular structure for a molecule the same as the arrangement of electron pairs? 3. If double bonds have two pairs of electrons why should they be treated as if they were single bonds when determining electron structure?

154 walk away with... Molecular structure describes how the atoms in a molecule are arranged in space Molecular structure can be predicted by using the valence shell electron pair repulsion (VSEPR)

155 51, 52, 59 12EOCs