Welcome to BondingOB: introduction to what bonding is, and

learning to draw (lots of) Lewis Dot Diagrams.

Chemistry is a course about what stuff is, and how this stuff reacts with other stuff. If a reaction happens, then bonds were formed, and most likely, other bonds were broken to let this happen. There is a lot to chemistry, the energy required or released, the properties of all that stuff, the conservation of energy, conservation of mass, and even conservation of CHARGE (coming soon).

Atoms of the elements are bonded together to make new stuff. There are a variety of bond types, all with rules and reasons. We’ll examine lots of them, but please realize that there is always way more going on - above our heads, and due to reasons we will not learn about in this course.

Types of bonds we will see:

1. Ionic - are bonds between cations and anions. The charges of the cations and anions will always net to zero.

2. Covalent - are bonds between two or more NON METALS. If there is a metal in the compound, it must be ionic. If there are no metals, then it’s covalent.

Types of bonds we will see:

3. Metallic - these are the connection thatmetal atoms make with each other when solid metals exist. They give rise to all the properties of metals. One type of metal atom bonds to itself.

4. Intermolecular - these are the kind of weak attractions between molecules. There are 3 different kinds of these, but two are nearly identical. I like these a lot.

Learning to draw Lewis Dot Diagrams. These will show us the outermost electrons, the VALENCE ELECTRONS, which are in the valence orbital. By reminding ourselves about these electrons it will help us to better understand what bonding can happen.

5. The outermost electrons are the VALENCE electrons6. The outermost electron orbital is the VALENCE ORBITAL.

7. Bonds always* form when atoms or ions end up with full outer orbitals, like the noble gases.

* of course there are exceptions, but not many, and we’ll get to these exceptions soon.

8.Dots will represent electrons.

9.Lewis dot diagrams will only show valence electrons, not the inside electrons. The inside electrons do not participate in the bonding anyway.

Electron Orbitals

10. The first orbital is tiny, it only holds 2 electrons at most.

11. The 2nd orbital is bigger, it can hold only up to 8 electrons (with a few exceptions!)

We won’t be drawing atoms bigger than 2 orbitals in our class, but they will be added in college chemistry, so be patient.

12. Together we’ll draw a few atoms, and ions, then YOU will continue these charts which will run from hydrogen to calcium.

Atom number Atom symbol Lewis Dot

(atom) Ion Symbol Lewis Dot(ion)

1 H H+1

2 He3 Li4 Be5 B

Atom number Atom symbol Lewis Dot

(atom) Ion Symbol Lewis Dot(ion)

1 H H· H+1 [H]+1

2 He He: --- ---3 Li :Li Li+1 [Li]+1

4 Be :Be: Be+2 [Be]+2

5 B :B Not in our class

Atom number Atom symbol Lewis Dot Ion Symbol Ion Dot

6 C7 N8 O9 F10 Ne11 Na

Atom number Atom symbol Lewis Dot Ion Symbol Ion Dot

6 C C Not in our class

7 N N N-3

8 O O O-2

9 F F F-1

10 Ne Ne --- ---11 Na Na Na+1 [Na]+1

Atom number Atom symbol Lewis Dot Ion Symbol Ion Dot

12 Mg13 Al14 Si15 P16 S17 Cl

Atom number Atom symbol Lewis Dot Ion Symbol Ion Dot

12 Mg Mg Mg+2 [Mg]+2

13 Al Al Al+3 [Al]+3

14 Si Si Not in our class

15 P P P-3

16 S S S-2

17 Cl Cl Cl-1

Atom number Atom symbol Lewis Dot Ion Symbol Ion Dot

18 Ar19 K20 Ca

Atom number Atom symbol Lewis Dot Ion Symbol Ion Dot

18 Ar Ar --- ---19 K K K+1 [K]+1

20 Ca Ca Ca+2 [Ca]+2

Lewis Dot diagrams for atoms show only valence electrons.

Lewis Dot diagrams for ions show the NEW Valence Orbital Arrangement.

Bonding class #2

OB: Metallic Bonds, More Lewis Dots, and the Octet Rule.

14. When sodium chloride forms from sodium metal and chlorine non metal, the atoms form ions first. To do this, the sodium TRANSFERS an electron to a chlorine atom .

15. The sodium becomes a sodium cation with a +1 charge

16. The chlorine becomes a chloride anion, with a -1 charge

17. Let’s draw the Lewis dot diagrams for the atoms, the ions, and then the compound.

ATOMS IONS COMPOUND

18. It’s important to note here, the sodium atom at 2-8-1 electron configuration becomes 2-8 as it loses one electron, becoming isoelectric to neon.

19. It loses enough electrons to get a perfect outer orbital, as defined by noble gases having the most perfect, or most stable electron orbitals of all.

20. The chlorine atom has a 2-8-7 configuration, gains one electron, and becomes 2-8-8, making it isoelectric to argon.

21. Both ions end up with perfect outer orbitals, both end up isoelectric to a noble gas.

22. Almost all ions follow the octet rule. 23. The octet rule is that when bonding all ions

will end up with eight outer most electrons, and when bonding, all non-metals bonding

together with other nonmetals in covalent bonds, will end up with 8 electrons in the

outermost orbitals.

24. This is a rule, but not a law. There are exceptions: some ions are too small, like Li.

Some atoms can squeeze 10 electrons, we love exceptions!

Compoundname

CompoundFormula

Cation Anion Lewis Dot Diagram

Magnesium oxide MgO Mg+2 O-2

LiF

CaCl2

Sodium… S-2

25. Fill in this chart!

Compoundname

CompoundFormula

Cation Anion Lewis Dot Diagram

Magnesium oxide MgO Mg+2 O-2

Lithium fluoride LiF Li+1 F-1

Calcium chloride CaCl2 Ca+2 Cl-1

Sodium… Na2S Na+1 S-2

Copy this table BIG, leave enough room for the dot diagrams!

Compoundname

CompoundFormula

Cation Anion Lewis Dot Diagram

Sodium… S-2

Cesium oxide

Compoundname

CompoundFormula

Cation Anion Lewis Dot Diagram

Sodium… Na2S Na+1 S-2

Cesium oxide Cs2O Cs+1 O-2

26. Why is the formula for aluminum oxide Al2O3 and not some other ratio?

Al O

O

Al

O

Each metal

atom is 2-8-3 and

needs to become

2-8 a +3

cation.Follow

the octet rule!

Each nonmetal atom is 2-6 and

needs to

become 2-8 a -2

anion.Follow

the octet rule!

Why is the formula for aluminum oxide Al2O3 and not some other ratio?

Al O

O

Al

O

A PERFECT TRANSFER OF ELECTRONS, 6 FROM Al + 6 INTO OXGYEN

27.Draw the UGLY Lewis dot diagram for

Magnesium Nitride Aluminum Oxide

28. This kind of bonding is to explain how metal atoms stick together to form solid metals. Literally, how does 6.02 x 1023 atoms of copper stick together so you can weigh 64 grams of copper on the scale? For the same reason that these atoms can stick together, nearly all of the properties of metals can be explained at the same time.As usual, it’s all about the electrons, where they are, what they’re doing, and how fast they can move.

First, let’s name a few properties of metals…

29. Metals are (you better learn the definitions of these ASAP)

Malleable Ductile Conduct electricity

Form cations Have higher densities than non metals

Have low Specific Heat Capacities than non metals etc.

These main properties can be explained by how we “understand” the metals to be bonded together.

Draw this diagram quickly…

30. Metals are understood to exist as packed cations, surrounded by loose valence electrons. These valence electrons can move quickly (near the speed of light) if they have to.

The positives balance the negatives since they are all atoms. Protons = electrons.

31. Imagine smashing the metal with a hammer to make the metal exhibit its malleable nature. The cations will be crushed closer together, and would repel, but the loose valence electrons flow to offset this excess positive charge.

Same when you squish it into a wire.

32. Imagine a flow of electrons (electricity) in from the left side. As electrons flow into the metal, there are too many negative electrons for the cations, so the excess electrons flow out the other side (the flow of electrons is electricity!).

33. The cations are awash in a sea of loose valence electrons.

Bonding Class #3OB: introduction to covalent bonding

34. Covalent bonding is when 2 or more nonmetals share their valence electrons to bond.

35. They do not transfer them like ionic compounds do.

36. With ionic bonding, there is a TRANSFER OF ELECTRONS FROM METAL → NONMETAL They still follow the octet rule (mostly). Ionic bonds require a metal to be first in the formula.

Ionic bonds make formula units (FU’s).

37. In Covalent Bonding, there is A SHARING OF THE VALENCE ELECTRONS,

AND THE ATOMS WILL FOLLOW THE

OCTET RULE. 38. NO METALS in any covalent bonds.39. Covalent bonds form molecules.

40. Molecules form with covalent bonds (sharing electrons)

by following the octet rule almost every time.

41. Let’s draw Lewis Dot Diagrams

H2

F2

41. Let’s draw Lewis Dot Diagrams

H2 H H

F2 Br Br

42. In covalent bonds, all atoms get to share enough electrons so

that they get full valence orbitals at least some of the time.

43. These bonds previous are all SINGLE BONDS because they only share a single pair of electrons (one electron from each atom).

44. They are also NONPOLAR bonds because there is NO DIFFERENCE IN electronegativity value between the atoms.

F2 + H2 have SINGLE NONPOLAR COVALENT

bonds

45. Draw the Lewis Dot Diagrams

HCl

H20

Let’s see if we can draw the Lewis Dot Diagrams for HCl and then, water.

H ClThe hydrogen atom in black has one valence electron. The chlorine, in red, has 7 valence electrons. Together the hydrogen gets to borrow one electron from chlorine to fill its tiny orbital, and chlorine gets to borrow one electron from hydrogen to fill its larger orbital (octet rule).

H O H

Here, the hydrogen are black again, and need to borrow one electron from oxygen each, to fill their tiny orbitals. Oxygen borrows one electron from each of the hydrogen atoms, to fill up its larger orbital (octet rule). Water is bent, don’t forget! We’ll learn why soon enough!

The red/black colors are not important, “just for seeing” it better as we learn at the beginning.

H Cl

H O H

The bond between H and Cl is between atoms with 2 different electronegativity values. What are their EN Values?

46. Is this a polar bond, or a nonpolar bond?

How about here? What are the EN Values for H and for O?

47. Is this a polar or nonpolar bonds?

H Cl

H O H

48. NAME THIS BOND

_____________________________

49. There are 2 identical bonds here (both H-O). Name THESE BONDSThere are 2…

___________________________

H Cl

H O H

48. NAME THIS BOND

SINGLE POLAR COVALENT

49. There are 2 identical bonds here (both H-O). Name THESE BONDSThere are 2…

SINGLE POLAR COVALENT

H Cl

H O H

50. Another way to draw this, with a lot less dots, is called a structural diagram. With a structural diagram, we only show the bonds, with short lines indicating shared electrons. A single dash represents a single covalent bond. Draw both of these molecules without dots, with structural diagrams.

H Cl H―Cl

H O H

Structural Diagrams

H―OH

This is a bit turned, but molecules move in 3 dimensions. It’s fine this way, or pointing in any other way.

51. Draw the Lewis Dot Diagram for AMMONIA (NH3), then the structural diagram. NAME THESE 3 BONDS TOO.

Think first:

N Nitrogen has 5 valence electrons, and they will be paired up in a Lewis dot diagram (and real life) because this is more stable.

To bond, one pair will have to open up to connect with3 hydrogen atoms.

NBring in the 3 hydrogen atoms… H

H H

NHH

H H―N―HH

Ammonia as Lewis Dots, and as a structural diagram.

Checking the electronegativity values, we see that H has a 2.2 while N has a 3.0

These bonds are all single polar covalent.

52. Draw Lewis Dot Diagram, and Structural Diagram for Methane, CH4

Determine exactly what types of bonds are present in this molecule.

Draw Lewis Dot Diagram, and Structural Diagram for Methane, CH4

Determine exactly what types of bonds are present in this molecule.

CH

HHH

H―C―HH

H

Electronegativity values of 2.2 for H, and2.6 for N, so there are

4 single polar covalent bonds in a molecule of CH4

53. The greater the difference in electronegativity values between two atoms, the greater the polarity of the bond. This works like little +/- magnets. Some magnets are stronger (greater EN difference) and some magnets are weaker (lesser EN difference). Fill in this chart, and then RANK from the greatest polarity of the bond (1), to the weakest (5).

Polarity rank

Molecule/name

EN #1

EN#2

EN diff Structural diagrams

H2hydrogen 2.2 2.2 0 H―H

PCl3

OF2

HBr

HI

Polarity rank

Molecule/name

EN #1

EN#2 EN diff Structural

diagrams

5 H2hydrogen 2.2 2.2

zerothis is a nonpolar

bondH―H

1PCl3

phosphoru

strichloride

2.2 3.2 1.0Cl―P―Cl

3 OF2

oxygen difluoride

3.4 4.0 0.6

2 HBr hydrogenbromide

2.2 3.0 0.8 H―Br

4 HI hydrogen

iodide2.2 2.7 0.5 H―I

F F

O

Cl

54. How do 2 oxygen atoms stay together in O2?Let’s draw two atoms Lewis Dot to start our thinking.

O O

55. How many electrons does EACH atom of oxygen need to complete the octet? Can they do this for each other?

Hint: move the bottom pairs of electrons to the open sides.

O O Squeeze them together now (this requires you

redraw)

O OIn order to both get an octet, the oxygen atoms must share 2 pairs of electrons with each other.

56. This gives the oxygen molecule a

DOUBLE COVALENT BOND.

57. Since each atom of oxygen has the same electronegativity value, it’s proper to call this a:

Double Non-Polar Covalent bond

O OO O

Bonding Class #4

OB: become masterful with both the Double and the Triple Covalent Bonds, plus some practice drawing structural diagrams for larger molecules

58. O2 Oxygen How does it bond? (review)

O OEach oxygen needs to gain 2 electrons to fill it’s valence

orbital. Each oxygen must lend 2 electrons to the other.

O OWill become…

O O Drawn structurally… this way: O=OEach O has a 3.4 electronegativity value, so this is a

Double Non-Polar Covalent Bond

59. Looking at the HONClBrIF twins, in order, let’s figure out the kinds of bonds that they all have… (draw dots or structural's to think)

H2

O2

N2

Cl2

Br2

I2

F2

Looking at the HONClBrIF twins, in order, let’s figure out the kinds of bonds that they all have… H2

O2

N2

Cl2

Br2

I2

F2

Single non-polar covalent H―H

Double non-polar covalent O=O

???

Single non-polar covalent Cl―Cl

Single non-polar covalent Br―Br

Single non-polar covalent I―I

Single non-polar covalent F―F

Time for some

Thinking!

N N60.

We’ll need two nitrogen atoms, which both need to follow the octet rule. How many electrons do they need to borrow from each other?

To do that, we’ll have to rearrange the electrons so that they can share them with each other.

N N Will shift →

And structurally this will become:

N N61. Nitrogen shares 3 pairs of electrons, it makes a triple nonpolar covalent bond

N N

62. Covalent bonds are between 2 or more nonmetals, and usually follow the octet rule.

63. Covalent bonds can be: SINGLE, DOUBLE, or TRIPLE

64. Covalent bonds can also be: POLAR or NON-POLAR

65. Let’s draw electron dot diagrams and then STRUCTURAL diagrams for these compounds.

DOTS Structural

C2H6

C2H4

Let’s draw electron dot diagrams and then STRUCTURAL diagrams for these compounds.

DOTS Structural

C2H6

C2H4

H―C―C―HH H

H H

C C H H

H HH H

C CH H

H H

C C H H

H H

66. Let’s draw electron dot diagrams and then STRUCTURAL diagrams for these compounds. Name each bond type.

DOTS StructuralC2H2

C3H8

Let’s draw electron dot diagrams and then STRUCTURAL diagrams for these compounds.

DOTS Structural

C2H2

C3H8

C CH H H―C C―H

C C C H H

H HH H

H

HH―C―C―C―H

H H H

H H H

H-C bond is single polar covalentC-C bond is triple nonpolar covalent

H-C bond is single polar covalentC-C bond is single nonpolar covalent

67. Let’s draw electron dot diagrams and then STRUCTURAL diagrams for these compounds. Name each bond type.

DOTS Structural

CO2

AsCl3

DOTS Structural

CO2

AsCl3

O=C=OC O OCarbon dioxide is a STRAIGHT molecule

Cl―As―Cl

Cl

As ClCl

Cl

O = C bond is a double polar covalent bond

As – Cl bond is a single polar covalent bond

Butane goes camping with some people in blue tanks, to run the stoves. It’s a repeating type molecule, a little chain really. The formula is C4H10. 68. Draw the Structural Diagrams now. 69. Name both kinds of bonds in this molecule

Butane formula is C4H10.

H―C―C―C―C―H H H H H

H H H H

The C―C bond is single nonpolar covalentand the C―H bond is single polar covalent

70. Draw Dot diagrams and structural diagrams for oxygen dibromide. 71. Name the bonds between Br + O.

Structurally, this becomes…

These bonds are SINGLE POLAR COVALENT bonds.

EN diff of 3.4 - 3.0 = 0.4

O Br

Br OBr

Br

72. Draw structural diagrams for carbon tetrachloride.

73. Name the bonds between carbon and chlorine.

CCl4 The Electronegativity difference between Cl - C is 3.2 – 2.6 = of C is 0.6, These are all single polar covalent bonds.

Cl C ClCl

ClCl―C―Cl

Cl

Cl

Bonding Class #5

OB: more practice with Lewis Dot diagrams, Structural Diagrams, and we get to meet the

weird hybrid bonds of ozone, carbon monoxide, PCl5, and NO2

Lots of paper, fresh minds, periodic tables at the ready!

74. First draw 2 diagrams, (structural or dots), and determine the types of bonds present.

for CaO and for CO2.

calcium oxide Carbon Dioxide CO2.

[Ca]+2[ O]-2 This is an IONIC bond, due to the transfer of electrons from calcium to oxygen.

This has two double polar covalent bonds, one to each side. The molecule is STRAIGHT due to these bonds.

75. Make sure you don’t forget that metals will only make IONIC bonds!

Now, let’s look over the model of NaCl crystal (and the diagram here).

Each Na+1 cation is surrounded by 6 Cl-1 ions.The reverse is also true, Each Cl-1 ion is surrounded by 6 Na+1 cations

76.We can say that the COORDINATION NUMBER for chloride is 6, and that the COORDINATION NUMBER for the sodium cations is also 6.

Big deal, right?77. Well, the coordination number, at the ion level, will give rise to specific shapes to the salt crystals when they grow up to be big enough to see with your eyes.

Uric Acid crystals look like this:

Cute little swords, really.They cause gout, which is what can happen to your teacher from time to time.

The last time my sister in law Andrea made Portobello Mushroom appetizers for Christmas Eve 3 years back, I woke up at 2 AM in pain and was unable drive home the next day, because of those crystals (all making a home in my big toe!).

78. Let’s work on carbon monoxide now, how does it bond together?CO2 is so important, straight line, two double polar covalent bonds, but what about it’s little cousin, CO?

C O

This will be an exceptional bond, with a cool name too. Here goes…

C OCan we share 4 + 6 to get two octets? This is tricky!

C OLooks like a polar double

bond is in order, but

carbon has no octet yet.

86

C OSo, oxygen will

“lend” 2 of it’s unshared electrons to the bonding “mix”, and it keeps an

octet, and although they are not bonded in the

same way, carbon “gets” an octet too.

C O79. It will form a double polar covalent bond and also form what’s called aCOORDINATE COVALENT BOND.The oxygen electrons coordinate this situation so that carbon “gets” an octet in a sort-of cheating way. Weird, but it happens!

C=O~ My shorthand for this type of bond

80. Phosphorus pentachloride is up next, a real but weirdo molecule…It’s used as part of fertilizer preparation, and other chemical reactions. Draw the dot diagram.

Phosphorus Pentachloride PCl5

P

Cl

Cl

Cl

ClCl

81. This compound breaks the octet rule, how many electrons does phosphorus end up with? How is this possible? Isn’t it a rule???

The compound called 2-pentene has five carbons in a chain, similar to octane (8) that we drew yesterday. The #2 means that the one double bond fits between the 2nd and 3rd carbon in the chain. Only hydrogen atoms are bonded to this set of five carbon atoms.

82. Can you draw the dots here, then the structural diagram? (I dare you to try, no talking 2 minutes! My kids can do anything!)

The compound called 2-pentene has five carbons in a chain, similar to octane (8) that we drew yesterday. The #2 means that the one double bond fits between the 2nd and 3rd carbon in the chain. Only hydrogen atoms are bonded to this set of five carbon atoms. Can you draw the dots, and the structural diagram? (I dare you to try, no talking 2 minutes! My kids can do anything!)

H―C ―C C ―C―C―HH

H

H H

HH

H HC5H1

0

John Adams High School, in Ozone Park, New York

City at left. I graduated in 1978 on that sidewalk out in front of the school. We

had 2 graduations that day, 975 kids were too

many for one ceremony!

Both my sister and my brother worked here, and they had the best crumb buns in the world. And the cookies were fab too.Those stairs go up to the A Train to Manhattan. It’s only the subway in Manhattan, most of Queens the subway runs on the el.

83. Ozone and Oxygen, both pure oxygen, but one you breathe to live, one you breathe to die. (sorry).

O3 vs. O284. You already know that oxygen has a double, nonpolar, covalent bond. No need to review (right?).

85. Ozone is an ALLOTROPE of oxygen. 86. Allotropes are pure forms of an element but due to different bonding, they have different properties. Other allotropes are carbon in the graphite mode and carbon in the diamond mode! 87. Try to bond 3 oxygen atoms…

O

O O

O

OO88. With ozone (and other molecules, like NO2, the electrons can’t add up to full octets all around. In this case, the oxygen atoms become “most” stable by making a double bond and a single bond, which RESONATES, back and forth.

It’s a resonating bond.

87

O

O O

O

OO89. In reality, this switching back and forth is constant, and becomes, 2 “one and a half bonds” all the time. Scientists know this because they can measure the bond lengths. Single bonds are longer than double bonds. These “resonating” bonds are 1½ sized all of the time. Getting the ozone bonding, the RESONATING bonding correct on the regents is a very personal thing for me. If you have any kindness in your heart for me, please remember this bond and how it is a “hybrid” bond (abnormal, like me!) (90)

Bonding Class #7

OBJECTIVE:91. Defining the 3 kinds of Intermolecular Bonding: the weak attractions between molecules, much weaker than ionic or covalent bonds, but they are important and have a real effect on the compounds

Quick review….

92. Ionic bonds form between metals (that lose electrons) and nonmetals (that gain electrons). The transfer of electrons result in the formation of neutral ionically bonded compounds, such as NaCl, MgO, or CuCl2

93. Covalent bonds form between 2 or more nonmetals (no metals ever) by sharing electrons. The molecules that form will have single, double or triple bonds, and atoms follow the octet rule. Examples are water, CH4, and CO2.

94. These bonds are all inside the compound.

95. There are three kinds of INTERMOLECULAR BONDS, bonds formed by the molecules with each other. These are all MUCH WEAKER that inside the compound bonds, but they are important.

96. Weakest to strongest they are: electron dispersion force, dipole interaction, and hydrogen bonding.

97. When I was in college there were only 2 kinds, electron dispersion and dipole interaction. Hydrogen bonding is very similar to dipole interaction, and we’ll see how they work today.

98. The weakest is the electron dispersion force. It’s created by the movement of electrons.

99. Electron Dispersion forces. Each of these F2 molecules has a 2-7 doubled electron configuration. Each atom has 9 electrons, the molecules have 18 electrons.

100. When these electrons all “move” to one side, for a nanosecond, there will be a temporary dipole created, a positive side, and a negative side of the molecule.

This allows for the weakest of temporary attractions to exist.

F2 is a gas at STP, because the kinetic energy at 273 Kelvin exceeds the attractive force of the electron dispersion forces, so it’s a

GAS.

Example one: fluorine F2

Electron Dispersion forces. 101. Each of these Cl2 molecules has a 2-8-7 doubled electron configuration. Each atom has 17 electrons, the molecules have 34 electrons.

When these electrons all “move” to one side, for a nanosecond, there will again be temporary dipoles. This happens more often than with fluorine, but not often enough to make a difference at 273 Kelvin.

This allows Cl2 to be a gas at STP, because the kinetic energy at 273 Kelvin exceeds the attractive forceof the electron dispersionforces, so chlorine is also a

GAS.

Example two: Chlorine Cl2

Electron Dispersion forces. 102. Each of these Br2 molecules has a 2-8-18-7 doubled electron configuration. Each atom has 35 electrons, the molecules have 70 electrons.

These electrons all “move” to one side, so many times per second that there will be a dipole created, a positive side, and a negative side of the molecule. This happens often enough that these attractive forces make Br2 a liquid!

The weak but constant intermolecular attractions accumulate.

The 273 Kelvin kinetic energy cannot overcome the intermolecular attractions, so bromine becomes a

liquid.

Example 3: Bromine Br2

Electron Dispersion forces. 103. Each of these I2 molecules has a 2-8-18-8-7 doubled electron configuration. Each atom has 53 electrons, the molecules have 106 electrons.

The electrons move so much, that a near constant dipole exists due to these electron dispersions.

This allows for the weakest of temporary attractions to exist at all times, which makes I2 a solid at STP.

The kinetic energy at 273 KelvinDOES NOT exceed the attractive force of the electron dispersionforces, so iodine is a

SOLID

Example 4: Iodine I2

104. The halogens clearly show how electron dispersion forces accumulate and then affect the molecules.

105. When there are dipoles, that means a positive and a negative side to a molecule (or a bond). Here, there are near permanent dipoles created by polar bonds but ONLY IN POLAR MOLECULES.

HH C H

H

S

Cl Cl

We have seen already that BONDS can be polar, due to either having a difference in electronegativity (they share electrons unevenly), or if they are IONIC and transfer electrons to bond (forming positive and negative ions)We’re about to start talking MOLECULAR POLARITY, is the whole molecule polar (different than bond polarity).

106. Molecular polarity is based upon SHAPE OF THE MOLECULE.

107. If the molecule is “balanced” it will be nonpolar.

108. The balance, or SYMMETRY we’re looking for is called RADIAL SYMMETRY

109. There are other symmetries, but they DO NOT matter.

110. In SCl2, the bonds are single polar covalent. The molecule itself is polar because it does not have radial symmetry. So, the sulfur will become positively charged most of the time, and the chlorine atoms will be negative most of the time.

S

Cl Cl

SCl2 had what sort of symmetry? Gingerbread

Man symmetry?

111. Methane molecules, which have polar bonds, has radial symmetry. It’s NON POLAR112. Radial Symmetry offsets bond polarity, the

molecule is nonpolar. SCl2 will be liquid at room temp while CH4 is a gas. Why???H

H C HH

S

Cl Cl

S

Cl Cl

S

Cl Cl

S

Cl ClS

Cl ClS

Cl Cl

113. Draw these: All the positive sulfur atoms are nearly permanently attracted to the negative chlorine atoms. The EN difference in a polar molecule can create intermolecular bonds called dipole attractions.

HH C H

H

HH C H

HHH C H

H

HH C H

H

114. Draw these. These methane molecules (nonpolar) have nearly no attraction to each other, so they will be gas at room temperature. Dipole attraction is way less powerful than ionic or even covalent bonding, but it can affect the phase of the compound. Nonpolar molecules are hardly attractive to each other.

Is there ANY attraction here between molecules?

115. Hydrogen bonding is exactly the same as dipole attraction, but, and it’s a SMALL but, hydrogen has to be present in the molecule.

116. H has a much smaller EN value than most other atoms, so when it’s included, like with water, the dipole it creates is usually much stronger than when it’s something like SCl2.

S

Cl ClO

H H

117. The EN difference between chlorine and sulfur is 3.2 – 2.6 = 0.6

118. The EN difference between oxygen and hydrogen is 3.4 – 2.2 = 1.2

119. This greater difference creates a “stronger” dipole. Strong enough that we now have to give it a new name. Instead of just calling it a strong dipole attraction, we call it hydrogen bonding.

O

H H

O

H H

O

H H

O

H H

O

H H

O

H H

120. Draw these now. All of the negative oxygen are magnetically attracted to the positive hydrogen atoms in nearby molecules. This is an intermolecular attraction. Hydrogen bonding is the strongest of the 3 intermolecular attractions.

121 Give an example molecule (or formula unit) for each type of bond:

IonicSingle nonpolar covalentSingle polar covalentDouble nonpolar covalentDouble polar covalentTriple non polar covalentTriple polar covalentCoordinate covalentResonantIonic + Covalent at the same timeBreaks the octet rule (more than 8e-)Breaks the octet rule (less than 8e-)

Give an example molecule (or formula unit) for each type of bond:

Ionic……………………………………………………………………………………..NaClSingle nonpolar covalent……………………………………………………F-FSingle polar covalent………………………………………………………….H-ClDouble nonpolar covalent………………………………………………….O=ODouble polar covalent………………………………………………………..O=C=O (both)Triple non polar covalent…………….…………………………………….NΞNTriple polar covalent………………………………………………………….NΞC-HCoordinate covalent…………………………………………………………..carbon monoxideResonant……………………………………………………………………………….ozone O3

Ionic + Covalent at the same time………………………………….CuSO4·5H2O*Breaks the octet rule (more than 8e-)………………………….PCl5Breaks the octet rule (less than 8e-)…………………….………H-H (too small)* Also has hydrogen bonding as well. (wow!)

Bonding Class #8OB: master relative oxidation numbers,

review all bonding for celebration tomorrow

A long, long time ago, in a galaxy, far, far away…

This is going to be great!

122. We learned about oxidation numbers, those little

positive and negative numbers in the corners of the periodic table, that told us what ratios of atoms to atoms molecular compounds make.Time to revisit them.

123. Hydrogen has a +1 and a -1 oxidation number.

Oxygen has only a -2 oxidation number.

To “make” molecules, you have to combine atoms to atoms, so that the sum of the oxidation number is zero. These are numbers, not ion charges!

Since oxygen is only a -2, it will take two +1 hydrogen atoms to make a molecule. That is why the formula is H2O, and that’s why H3O or HO is not a real compound.

124. Let’s determine the relative oxidation numbers in these molecules…

HCl

CH4

CO2

Let’s determine the relative oxidation numbers of the atoms in these molecules…

HCl H+1 Cl-1 (+1) + (-1) = 0

CH4 C-4 H+1 (-4) + 4x(+1) = 0

CO2 C+4 O-2 (+4) + 2x(-2) = 0

Sulfur dioxide SO2 S+4 O-2 O-2 (0)Chromate ion CrO4

-2 Cr+6 O-2 O-2 O-2 O-2 (-2)125 Permanganate ion126 NH3

127 NaOH128 KClO3

129 Carbon monoxide130 Carbon dioxide131 Dihydrogen sulfate132 Nitrate ion133 Nitrogen dioxide

134 Phosphorus trichloride

Sulfur dioxide SO2 S+4 O-2 O-2 (0)Chromate ion CrO4

-2 Cr+6 O-2 O-2 O-2 O-2 (-2)Permanganate ion MnO4

-2 Mn+6 O-2 O-2 O-2 O-2 (-2)ammonia NH3 N-3 H+1 H+1 H+1 (0)

Sodium hydroxide NaOH Na+1 O-2 H+1 (0)Potassium chlorate KClO3 K+1 Cl-5 O-2 O-2 O-2 (0)Carbon monoxide CO C+2 O-2 (0)

Carbon dioxide CO2 C+4 O-2 O-2 (0)Dihydrogen sulfate H2SO4 H+1 H+1 S+6 O-2 O-2 O-2 O-2 (0)

Nitrate ion NO3-1 N+5 O-2 O-2 O-2 (-1)

Nitrogen dioxide NO2 N+4 O-2 O-2 (0)Phosphorus trichloride PCl3 P+3 Cl-1 Cl-1 Cl-1 (0)

Intermolecular bonding system Jeopardy!

135. It keeps ammonia NH3 together as a liquid

what is…

136. It keeps Br2 bromine a liquid, but iodine I2 a solid

137. It keeps phosphorus trichloride PCl3 together as a liquid

Intermolecular bonding system Jeopardy!It keeps ammonia NH3 together as a liquid What is hydrogen bonding?

It keeps Br2 bromine a liquid, but iodine I2 a solid What is the electron dispersion force or electron dispersion attraction?

It keeps phosphorus trichloride PCl3 together as a liquid What is the dipole attraction force?

138. In one sentence explain the difference between bond polarity and molecular polarity. Who has the guts to stand and orate this one?

Bond polarity is when there is a difference in electronegativity value between two atoms that are bonding. All ionic bonds are polar, but for covalent bonds we have to check table S.

Molecular polarity has to do with molecular shape.If a molecule has radial symmetry, it is a nonpolar molecule.A molecule that doesn’t exhibit radial symmetry is polar.A polar molecule water A non polar molecule CCl4

In Queens, especially in Ozone Park, you can get on the A train and go to Brooklyn. Then you get off, cross the platform, and go back to Ozone Park in Queens. You can do this over and over all day long, all night long, all for one price.

You can resonate back and forth from Queens to Brooklyn.

The bonds in ozone O3 resonate back and forth, they are exceptional bonds, but they exist.

In reality the bonds are “both” 1½ sized rather than small doubles and bigger single bonds.

139. Once and for all, with a little dot diagram, and one sentence, explain how carbon monoxide bonds together.

C O

Once and for all, with a little dot diagram, and one sentence, explain how carbon monoxide bonds together.

C OIt’s called a double polar covalent bond (the bottom 2 pairs of electrons) and a coordinate covalent bond, which means oxygen just “lends” 2 electrons into the mix so carbon “gets” an octet too.

140. True or False?

1. Ionic bonds can be double or single bonds

2. Covalent bonds cannot be nonpolar bonds

3. Oxygen molecules have double polar covalent bonds

4. Nitrogen molecules have double nonpolar covalent bonds

5. Hydrogen atoms can make single or double covalent bonds

6. Oxygen atoms must make double bonds ONLY

7. Water is sometimes a straight line molecule by shape

8. Molecules with polar bonds can never be non polar molecules

9. Molecules with nonpolar bonds only can never be polar molecules

10. The weakest intermolecular bond is the dipole force of attraction

True or False? ALL FALSE!!!

1. Ionic bonds can be double or single bonds No, just magnetic ionic

2. Covalent bonds cannot be nonpolar bonds No, F2 or Cl2 are nonpolar

3. Oxygen molecules have double polar covalent bonds No, double nonpolar

4. Nitrogen molecules have double nonpolar covalent bonds No, triple nonpolar

5. Hydrogen atoms can make single or double covalent bonds No, only single

6. Oxygen atoms must make double bonds ONLY No, in water they make 2 singles

7. Water is sometimes a straight line molecule by shape No, always, always bent!

8. Molecules with polar bonds can never be non polar molecules No, CO2 or CH4

9. Molecules with nonpolar bonds only can never be polar molecules No, NBr3

10. The weakest intermolecular bond is the dipole force of attraction No, electron dispersion forces are weakest, watch them in Group 17

Study tonight, and every night.Do the 2 drills in BONDING section on homepage.

Monday we celebrate

20 multiple choice from old regents exams, then

Draw a few Lewis dot diagrams (also called electron dot diagrams)

+ 2 word answer problems