Welcome to Bonding OB: introduction to what bonding is, and learning to draw (lots of) Lewis Dot Diagrams. 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. Electrons are TRANSFERRED. 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 its Types of bonds we will see: 3. Metallic - these are the
connection that metal 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 Learning to draw Lewis Dot Diagrams. These will show us the outermost electrons, the VALENCE ELECTRONS, which are in the valence orbital. 5. The outermost electrons are the VALENCE electrons 6. 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, and well get to these exceptions soon.
8. The 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 in high school. Electron Orbitals 10. The first orbital is tiny, it only holds 2 electrons at most. 11. The 2nd orbital is bigger, it fills up with 8 electrons (with a few exceptions!) 12. Together well draw a few atoms, and ions, then YOU will continue these charts which will run from hydrogen to calcium. Atom number Atom symbol 1
H 2 He 3 Li 4 Be 5 B Lewis Dot
(atom) Ion Symbol H Lewis Dot (ion) +1 X X Atom number Atom symbol
Lewis Dot (atom) Ion Symbol Lewis Dot (ion) 1 H H H+1 [H]+1 2
He He X X 3 Li Li Li+1 [Li]+1 4
Be Be Be+2 [Be]+2 5 B B X X Atom number
Atom symbol 6 C 7 N 8 O 9 F 10
Ne 11 Na Lewis Dot Ion Symbol Ion Dot X X X X
Atom number Atom symbol Lewis Dot Ion Symbol Ion Dot 6 C C X
X 7 N N N-3 8 O O O-2 9
F F F-1 10 Ne Ne X X 11 Na
Na Na+1 [Na]+1 Atom number Atom symbol 12 Mg 13 Al 14
Si 15 P 16 S 17 Cl Lewis Dot Ion Symbol Ion Dot
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 X
X 15 P P P-3 [P] 16 S S S-2
[ S ]-2 17 Cl Cl Cl-1 [ Cl ]-1 -3 Atom number Atom symbol 18
Ar 19 K 20 Ca Lewis Dot Ion Symbol Ion Dot x x
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. 20. 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 . 21. The sodium becomes a sodium cation with a +1 charge 22. The chlorine becomes a chloride anion, with a -1 charge 23. Lets draw the Lewis dot diagrams for the atoms, the ions, and then the compound. ATOMS ONLY
IONS ONLY COMPOUND 23. Lets draw the Lewis dot diagrams for the atoms, the ions, and then the compound. SEPARATE FAR APART ATOMS ONLY IONS ONLY SNUGGLY COMPOUND 24. NOTE: the sodium atom at 2-8-1 electron configuration
becomes 2-8-0 (just 2-8 and an EMPTY valence orbital) as it loses one electron, becoming isoelectric to neon. 25. Sodium loses enough electrons to get a perfect outer orbital, as defined by the noble gases. They all have perfect, and the most stable electron orbitals of all. 26. The chlorine atom has a 2-8-7 configuration, gains one electron, and becomes chloride anion with 2-8-8, making it isoelectric to argon. 27. Both ions end up with perfect outer (valence) orbitals, both ions end up isoelectric to noble gases. 28. Almost all ions follow the octet rule. 29. The octet rule is that all ions end up with eight outer most or valence electrons. When bonding, all non-metals bonding together with other nonmetals in covalent bonds, also end up with 8 electrons in the
outermost orbitals. 30. This is a rule, but not a law. There are exceptions: some ions are too small for 8 electrons like Li, or atoms too small, like H. There are some exceptions! 31. Fill in this chart! Compound name Compound Formula Cation Anion Magnesium oxide
MgO Mg+2 O-2 LiF CaCl2 Lewis Dot Diagram Copy this table BIG, leave enough room for the dot diagrams! Compound name Compound Formula
Cation Anion Magnesium oxide MgO Mg+2 O-2 Lithium fluoride LiF Li+1
F-1 Calcium chloride CaCl2 Ca+2 Cl-1 Lewis Dot Diagram Compound name Sodium Cesium oxide
Compound Formula Cation Anion S-2 Lewis Dot Diagram Compound name Compound Formula Cation
Anion Sodium Sulfide Na2S Na+1 S-2 Cesium oxide Cs2O Cs+1 O-2
Lewis Dot Diagram 32. Why is the formula for aluminum oxide Al2O3 and not some other ratio? Each metal atom is 2-8-3 and needs to become 2-8 a +3 cation. Follow the octet rule! Al
O O Al O Each nonmet al 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 33. Draw the UGLY Lewis dot diagram for Magnesium Nitride Aluminum Oxide
33. Draw the UGLY Lewis dot diagram for Magnesium Nitride Aluminum Oxide Metallic Bonding 34. First, lets name a few properties of metals Metallic Bonding 28. First, lets name a few properties of metals Metals are (you better learn the definitions of these ASAP) Malleable Ductile Conduct electricity Form into cations
Have higher densities than non metals Have lower Specific Heat Capacities than non metals These main properties can be explained by how we understand the metals to be bonded together. 35. Metals are understood to be: PACKED CATIONS SURROUNDED BY A SEA OF LOOSE VALENCE ELECTRONS. LOOK at this diagram 36. Metals are made up of NEUTRAL ATOMS. A hunk of metal is neutral. The positive protons balance the negative electrons as in all atoms. # Protons = # electrons. When the electrons become loose this neutrality remains, but the electrons can move to offset too much positive charge formed when the cations are smushed together.
37. 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 from each other causing a crack in the metal. The loose valence electrons flow to offset this excess positive charge. Same when you squish it into a wire. 38. Imagine a flow of electrons (electricity) in from the left side. As electrons flow into the metal, there are too many electrons for the cations, so the excess electrons flow out the other side (the flow of electrons is electricity!). 39. In metals, the cations are awash in a sea of loose valence electrons.
Bonding Class #3 OB: introduction to covalent bonding 40. Covalent bonding is when 2 or more nonmetals share their valence electrons to bond. 41. They do not transfer electrons, like ionic compounds do. 42. 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 (FUs). 43. In Covalent Bonding, there is A SHARING OF VALENCE ELECTRONS, AND THE ATOMS WILL FOLLOW THE
OCTET RULE. 38. NO METALS in any covalent bonds. 46. Molecules form with covalent bonds (sharing electrons) by following the octet rule almost every time. 47. Lets draw Lewis Dot Diagrams H2 F2 47. Lets draw Lewis Dot Diagrams H2 HH
F2 F F 48. In covalent bonds, all atoms get to share enough electrons so that they get full valence orbitals at least some of the time. 49. These bonds previous are all SINGLE NON POLAR COVALENT BONDS because they only share one pair of electrons (one electron from each atom) AND because there is no difference between the electronegativity values (Table S, both are 2.2 in H 2, or both are 4.0 in F2) H2 + F2 both have SINGLE NONPOLAR COVALENT bonds 50. 51. Draw the Lewis Dot Diagram for HCl and name the bond present.
HCl 51. Draw the Lewis Dot Diagram for HCl and name the bond present. H Cl Chlorine gets to borrow one electron from hydrogen to fill its larger orbital (octet rule). Hydrogen atoms cant get 8 electrons, they are too small, but it still follows the fill up my baby sized orbital when I bond rule. This is a single polar covalent bond. 51. Draw the Lewis Dot Diagram for H2O and name the bond present. H2O 52. Draw the Lewis Dot Diagram for H2O and name the bond present.
H2O 52. Draw the Lewis Dot Diagram for H2O and name the bond present. HO H Oxygen borrows one electron from each of the hydrogen atoms, to fill up its larger orbital (octet rule). Hydrogen atoms borrow the one electron they need to fill up the baby sized orbital too. Water is bent, dont forget! Well learn why soon enough! The bonds are both SINGLE POLAR COVALENT The red/black colors are not important, just for seeing it better as we learn at the beginning.
H Cl HO H 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. 53. Draw both of these molecules without dots, with structural diagrams. H Cl HO
H Structural Diagrams HCl O H H This is a bit turned, but molecules move in 3 dimensions. Its fine this way, or pointing in any other way. 54. Draw the Lewis Dot Diagram for AMMONIA (NH3), AND the structural diagram. NAME THE BONDS TOO. Think first:
N 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 with 3 hydrogen atoms. Bring in the 3 hydrogen atoms H H H H NH H HNH
H 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. 55. 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. H HCH
H H HCH H Electronegativity values of 2.2 for H, and 2.6 for N, so there are 4 single polar covalent bonds in a molecule of CH4 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). 56. Fill in this chart, and then RANK from the greatest polarity of the bond (1),to the weakest (5). Molecule formula + name
EN #1 EN #2 EN diff H2 hydrogen 2.2 2.2 0 PCl3
OF2 HBr HI Polarity rank Structural diagrams HH Molecule formula + name EN #1 EN #2
EN diff Polarity rank Structural diagrams H2 hydrogen 2.2 2.2 0 0
HH PCl3 phosphoru s trichloride 2.2 3.2 1.0 1 ClPCl Cl O OF2
oxygen difluoride 3.4 4.0 0.6 3 F HBr hydrogen bromide 2.2 3.0
0.8 2 HBr HI hydrogen iodide 2.2 2.7 0.5 4 HI
F 57. Draw two Lewis Dot diagrams of oxygen atoms to start our thinking. O O 58. How many electrons does EACH atom of oxygen need to complete the octet? Can they do this for each other? O O 59. The atoms share 2 pairs of electrons, making
a. OO DOUBLE NONPOLAR COVALENT BOND. Why is this a NONPOLAR bond? _____ Bonding Class #4 OB: become masterful with both the Double and the Triple Covalent Bonds, plus some practice drawing structural diagrams for larger molecules 59. Looking at these HONClBrIF twins, in order, lets figure out the kinds of bonds that they all have (draw structural diagrams) H2 O2 Cl2 Br2 I2
F2 59. Looking at these HONClBrIF twins, in order, lets figure out the kinds of bonds that they all have (draw structural diagrams) H2 Single non-polar covalent HH O2 Double non-polar covalent O=O Cl2 Single non-polar covalent ClCl Br2 Single non-polar covalent
BrBr I2 F2 Single non-polar covalent II Single non-polar covalent FF 61. Draw a Lewis Dot Diagram for a nitrogen atom 62. Draw a molecule of nitrogen
in this box How many electrons does each atom need to meet the octet rule? Draw a Lewis Dot Diagram for another nitrogen atom 61. Draw a Lewis Dot Diagram for a nitrogen atom How many electrons does each atom need to meet the octet rule? Draw a Lewis Dot Diagram for another
nitrogen atom N Each nitrogen atom needs to gain 3 electrons for an octet. N 62. Draw a molecule of nitrogen in this box NN N N
Lewis Dot Diagram and structural diagram. This is a triple nonpolar covalent bond. 63. Nitrogen makes a triple nonpolar covalent bond because Each atom of nitrogen needs to borrow 3 electrons from the other, so they share three pairs of electrons (thats the triple part) Since each Nitrogen atom has an electronegativity of 3.0, there is NO DIFFERENCE in electronegativity values, so the bond is NONPOLAR. 64. Lets draw electron dot diagrams and then STRUCTURAL diagrams for these compounds. The first one is called ETHANE, with a formula of: DOTS C2H6 Structural
64. Lets draw electron dot diagrams and then STRUCTURAL diagrams for these compounds. The first one is called ETHANE, with a formula of: C2H6 DOTS H H H C CH H H C-C bonds are single nonpolar covalent C-H bonds are single polar covalent
Structural H H HCCH H H 65. Lets draw electron dot diagrams and then STRUCTURAL diagrams for these compounds. This one is called ETHENE, with a formula of: DOTS C2H4 Structural 65. Lets draw electron dot diagrams and then STRUCTURAL diagrams for these compounds. This one is called ETHENE, with a formula of: C2H4
DOTS H H H H C C H Structural H C=C bonds are double nonpolar covalent C-H bonds are single polar covalent C C H H
66. Lets draw electron dot diagrams and then STRUCTURAL diagrams for these compounds. This is ETHYNE, with a formula of: DOTS C2H2 Structural 66. Lets draw electron dot diagrams and then STRUCTURAL diagrams for these compounds. This is ETHYNE, with a formula of: DOTS H C2H2 C CH CC bonds are triple nonpolar covalent C-H bonds are single polar covalent
Structural HC CH 67. Lets draw electron dot diagrams and then STRUCTURAL diagrams for these compounds. This one is called PROPANE, with a formula of: C3H8 DOTS Structural 67. Lets draw electron dot diagrams and then STRUCTURAL diagrams for these compounds. This one is called PROPANE, with a formula of: C3H8
DOTS H H H Structural C C CH H H H H H H
HCCCH H H-C bond is single polar covalent C-C bond is single nonpolar covalent H H H 68. Lets draw electron dot diagrams and then STRUCTURAL diagrams for these compounds. This one is called CARBON DIOXIDE, with a formula of: CO2 DOTS Structural 68. Lets draw electron dot diagrams and then STRUCTURAL diagrams for
these compounds. This one is called CARBON DIOXIDE, with a formula of: CO2 DOTS O C O Structural O=C=O O = C bond is a double polar covalent bond Carbon dioxide is a STRAIGHT molecule 69. Lets draw electron dot diagrams and then STRUCTURAL diagrams for these compounds. This one is called ARSENIC TRICHLORIDE, with a formula of:
AsCl3 DOTS Structural 69. Lets draw electron dot diagrams and then STRUCTURAL diagrams for these compounds. This one is called ARSENIC TRICHLORIDE, with a formula of: AsCl3 DOTS Cl As Cl Cl Structural ClAsCl
Cl As Cl bond is a single polar covalent bond 70. Lets draw electron dot diagrams and then STRUCTURAL diagrams for these compounds. This one is called BUTANE, with a formula of: C4H10 DOTS Structural 70. Lets draw electron dot diagrams and then STRUCTURAL diagrams for these compounds. This one is called BUTANE, with a formula of: C4H10
DOTS At some point, and I have reached this point, I promise to remember always that carbon MUST make 4 bonds, and that hydrogen can only make 1 bond. The dots are for thinking, I have thought about this enough, I wont draw the dots anymore because I am smart and I promise myself to remember this. If I forget how to count to 4 I will make a boo boo and I wont belly ache about it. It would be my fault, and Structural
H H H H HCCCCH H H H H The CC bond is single nonpolar covalent and the CH bond is single polar covalent 71. Lets draw electron dot diagrams and then STRUCTURAL diagrams for these compounds. This one is called OXYGEN DIBROMIDE, with a formula of: OBr2 DOTS Structural 71. Lets draw electron dot diagrams and then STRUCTURAL diagrams for
these compounds. This one is called OXYGEN DIBROMIDE, with a formula of: OBr2 DOTS O Br Br Structural O Br Br These bonds are SINGLE POLAR COVALENT bonds. Electronegativity difference: 3.4 - 3.0 = 0.4 72. Lets draw electron dot diagrams and then STRUCTURAL diagrams for
these compounds. This is called CARBON TETRACHLORIDE, with a formula of: CCl4 DOTS Structural 72. Lets draw electron dot diagrams and then STRUCTURAL diagrams for these compounds. This is called CARBON TETRACHLORIDE, with a formula of: CCl4 Cl DOTS Cl
Cl C Cl Cl Structural ClCCl Cl CCl4 The Electronegativity difference between Cl - C is 3.2 2.6 = of C is 0.6, These are FOUR single polar covalent bonds. 73. Draw a Lewis Dot diagram for CaO calcium oxide, and tell what sort of bond or bonds are present. CaO 73. Draw a Lewis Dot diagram for CaO calcium oxide, and tell what sort of bond or bonds are present.
CaO [Ca] +2 [ O] -2 These are IONIC BONDS. They are NOT double bonds. The positive charge (+2) adds to the negative charge (-2) and must sum to zero. Bonding Class #5 OB: Random Vocabulary, more practice with Lewis Dot diagrams, Structural Diagrams, and we get to meet the weird hybrid bonds of
ozone, carbon monoxide, PCl5, and NO2 Fresh minds, periodic tables at the ready! ALLOYS 74. Alloys are MIXTURES of either 2 or more metals, OR metals + nonmetals. They are not chemically bonded! They get mixed most often by melting them together since metals are solids and dont mix well otherwise. The resulting stuff is not a new substance, its a mixture of the original substances. The alloy has different properties than the original substances because they pack together mixed up. Most alloys are made for strength, non-corrosiveness, or beauty. 75. Examples of ALLOYS:
Sterling Silver made from SILVER + COPPER for strength Cast Iron (IRON + CARBON) for strength and non-corrosiveness Stainless steel (Fe + Cr) for strength and non-corrosiveness Brass (ZINC + COPPER) for durability and beauty and tubas. 76. Coordination numbers in ionic solids. Metals will only make IONIC bonds with nonmetals, like NaCl. 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 Sodium has a coordination number of 6. Chloride also has a coordination number of 6. 77. Since both are 6, we end up with a nice boxy shape. How the ions pack together, due to their sizes, and their relative charges (+1 to -1, or possibly +3 to -1, etc.) will set their shape when they bond millions of
ions together. How they fit as ions is how they look when the form into crystals big enough to see with your eyes. 78. Calcium carbonate CaCO3 has a much different shape than NaCl because these compounds have different coordination numbers The different ions fit together differently and they grow up to look different than the NaCl crystals. 79. Coordination number the number of the different kind of ions surrounding one kind of ion in an ionic solid
structure. How many Cl-1 surround each Na+1 in NaCl(S) or How many Na+1 surround each Cl-1 in NaCl(S) Big deal, right? 80. 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 Little swords, really. They cause gout, which is what can happen to your teacher from time to time. That is not MY foot, but my foot has looked like that on some days.
I am not sure what the coordination number for uric acid crystals is, but I can assure you it is not 6 and 6 like NaCl. Whatever the coordination number for little swords of pain are, that coordination number is the reason for a lot of pain in the world. Gout is more prevalent in men, but my Grandmother used to get it. She gave it to my father and his brother. My Aunt didnt get it but both of her sons have it. My good friend Carl gets it, but he didnt catch it from me, his came from his own father. 81. Lets work on carbon monoxide now, how does it bond together? CO2 is so important, straight line, two double polar covalent bonds, but what about its little cousin, CO? C O
CO So, oxygen can lend 2 of its unshared electrons to the bonding mix, and it keeps an octet, and carbon gets an octet too. 82. CO forms CO ~ C=O a double polar covalent bond and also forms whats called a COORDINATE COVALENT BOND.
The oxygen electrons coordinate this situation so that carbon gets an octet in a sort-of cheating way. Weird, but it happens! 81. My shorthand for this type of bond 84. Phosphorus pentachloride is up next, a real but weirdo molecule Its used as part of fertilizer preparation, and other chemical reactions. Common in chemistry world, new to you. Draw the dot diagram. Phosphorus Pentachloride Cl Cl
Cl P PCl5 Cl Cl 85. This breaks the octet rule by phosphorous ending up with TEN valence electrons, which normally is 2 too many! Its an exception to this rule. 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. In Queens this train runs on elevated tracks (called the EL). In Brooklyn and Manhattan this train runs underground (the subway) 86. Ozone and Oxygen, both pure oxygen, but one you breathe to live, one you breathe to die. (sorry). O3 vs. O2 You already know that oxygen has a double, nonpolar, covalent bond. No need to review (right?).
87. Ozone is an ALLOTROPE of oxygen. 88. 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! Try to bond 3 oxygen atoms 89 O O O O O O
90. With ozone (and some other molecules, like NO2, the electrons cant get 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. Its a resonating bond. O O 91. O O
O O In reality, this switching back and forth is constant, and becomes, two 1 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 1 sized all of the time. Bonding Class #6 92. 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. 93. 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 94. 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. Quick review.
94. Metallic bonds keep hunks of metal together, but are not really bonds, they are a way to explain how metals seem to stay stuck and how metals have certain important properties. 95. These bonds are all inside the compound, or for metallic bonds, are just weird. 96. 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. 97. Weakest to strongest they are: electron dispersion force, dipole interaction, and
hydrogen bonding. 98. The weakest is the electron dispersion force. Its created by the constant movement of the electrons in atoms, or compounds. Sometimes its called Electron Dispersion Attraction. 99 Example one: fluorine F2 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 at STP, the kinetic energy exceeds the attractive force of the electron dispersion attraction, so its a GAS. Electron Dispersion forces. Example two: Chlorine Cl2 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 stronger with chlorine, but not often enough to make a difference at 273 Kelvin. 102. Cl2 to be a gas at STP, because the kinetic energy at 273 Kelvin exceeds the attractive force of the electron dispersion forces, so chlorine is also a GAS. Electron Dispersion forces. Example 3: Bromine Br2 103.
Each of these Br2 molecules has a 2-8-18-7 doubled electron configuration. Each atom has 35 electrons, the bromine molecules contain 70 electrons! These electrons all move to about, so when a dipole is temporarily created, its stronger than with F2 or Cl2. At STP, this attractive force is stronger than the kinetic energy at STP. It makes Br2 a liquid! The weak but constant intermolecular attractions accumulate, and impact the phase of the substance. 104. The 273 Kelvin kinetic energy cannot overcome the intermolecular attractions, so bromine become is a
Electron Dispersion forces. Example 4: Iodine I2 105. 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 strong dipole forms over and over, making these molecules have many moments of attraction. This allows for the weakest of temporary attractions to exist so often that it makes I2 a solid at STP. 106. The kinetic energy at 273 Kelvin
DOES NOT exceed the attractive force of the electron dispersion forces, so iodine is a SOLID 107. At STP: The halogens, of group 17 are gases (fluorine and chlorine) or liquid (bromine) or solid (iodine) at STP. 108. This difference in phases is only due to the amount of intermolecular attraction caused by the movement of electrons in the molecules, and it is called ELECTRON DISPERSION FORCE and it is cased by the moment to moment dispersion of the electrons. dipole attraction 109. A dipole occurs when a one molecule ends up with a positive side and a negative side because of nonradial symmetry. Polar bonds can
add to the polarity as well. Here, there are near permanent dipoles created by polar bonds but ONLY IN POLAR MOLECULES. Draw Dipole Arrows here. Do these molecules have radial symmetry? S Cl H H C H Cl H dipole attraction 110. A dipole arrow quickly which side of a bond is + and which side gets the electrons in the bond more often ().
The dipole arrows DO NOT replace the bond dashes. Draw Dipole Arrows here. Do these molecules have radial symmetry? S Cl H H C H Cl H 111. Molecular polarity is based upon SHAPE OF THE MOLECULE. 112. If the molecule has RADIAL SYMMETRY, it is balanced and it will be nonpolar.
113. The balance, or SYMMETRY were looking for is called RADIAL SYMMETRY 114. There are other symmetries, but they DO NOT matter in chem. 115. 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? 116. Methane, which has polar bonds too, but also has radial symmetry. 117. Radial symmetry offsets that polarity, and the molecule is nonpolar. SCl2 will be liquid at room temp. while methane would be a gas. Why??? H S
Cl H C H Cl H 117. Radial symmetry offsets that polarity, and the molecule is nonpolar. SCl2 will be liquid at room temp. while methane would be a gas. Why??? Radial symmetry means that the polar bonds are balanced, so nonpolar molecules do not attract together. Methane molecules are all positive on the outside. Molecules that approach each other repel apart. Molecules like SCl2 are polar, the negative chlorine atoms of one molecule are attracted to the positive sulfur atoms in the next molecule. This is a dipole attraction that is nearly constant. Methane is a gas, sulfur dichloride would be a liquid at STP.
118. 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. S S Cl Cl Cl S Cl Cl S Cl S
Cl Cl Cl Cl 119. 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. H
H H C H H H H C H H H C H H H C H H H Is there ANY attraction here between molecules? Hydrogen Bonding 120. Hydrogen bonding is exactly the same as dipole
attraction, but, hydrogen has to be present in the molecule. 121. H has a much smaller electronegativity value than most other atoms, so when its included, like with water, the dipole it creates is usually much stronger than when its something like SCl 2. 1 atom 1 atom compound water SCl2 difference
electronegativity electronegativity H 2.2 O 3.4 1.2 S 2.6 Cl 3.2 0.6 The dipole of water is so much stronger than the dipole of sulfur dichloride because of the great difference in the electronegativity values. Its so big, it gets a new name.
122. 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 S Cl Cl hydrogen bonding. O H H
Electronegativity values and differences S___ Cl ___ difference _______ H___ O ___ difference _______ This super duper dipole has a new name: its HYDROGEN BONDING 123. Draw these now. All of the negative O H O H H
H 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. O H O H H O H H
O H H H 124 Give an example that contains each of these types of bonds bond type Ionic Single nonpolar covalent Single polar covalent Double nonpolar covalent Double polar covalent Triple non polar covalent Triple polar covalent Coordinate covalent
Resonant Ionic + Covalent at the same time Breaks the octet rule (more than 8e-) Breaks the octet rule (less than 8e-) example formulas 124 Give an example that contains each of these types of bonds bond type example formulas Ionic NaCl, CaO, KBr, MgO Single nonpolar covalent
H2 , F2, Cl2, Br2 Single polar covalent HCl, HF, HBr, HI Double nonpolar covalent O2 Double polar covalent CO2 (Each C to O bond) Triple non polar covalent N2 Triple polar covalent
NCH (N to C is triple) Coordinate covalent CO (double polar cov, PLUS) Resonant O3 (ozone park!) Ionic + Covalent at the same time CuSO45H2O (also has hydrogen bonding!) Breaks the octet rule (more than 8e-) PCl5 Breaks the octet rule (less than 8e-)
H2 or LiF (the lithium) Bonding Class #7 OB: master relative oxidation numbers, review all bonding for celebration tomorrow ----------------- CuSO45H2O* A long, long time ago, in a galaxy, far, far away 125. We learned about oxidation numbers, those little positive and negative numbers in the corners of the NON METALS on the periodic table, that help us decide what ratios of atoms to atoms that
molecular compounds make. Time to revisit them. 126. 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 thats why H3O or HO cannot form. 127. Lets determine the relative oxidation numbers in these molecules
HCl CH4 CO2 Lets determine the relative oxidation numbers of the atoms in these molecules HCl H CO2 C +1 +4 Cl O
-2 -1 (+1) + (-1) = 0 (+4) + 2x(-2) = 0 AsCl3 As+3 Cl-1 (+3) + 3x(-1) = 0 129 Sulfur dioxide SO2 Chromate ion CrO4-2
Permanganate ion 130 NH3 131 NaOH 132 KClO3 133 Carbon monoxide 134
Carbon dioxide 135 Dihydrogen sulfate 136 Nitrate ion 137 Nitrogen dioxide 138 Phosphorus trichloride S+4 O-2 O-2 (sums to: 0) Cr+6 O-2 O-2 O-2 O-2 (sums to: -2)
Sulfur dioxide SO2 S+4 O-2 O-2 ( sums to = 0) Chromate ion CrO4-2 Cr+6 O-2 O-2 O-2 O-2 (sums to -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 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 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 Na+1 O-2 C+4 P+3 Cl-1 H+1 (0) O-2 O-2 (0) Cl-1 Cl-1 (0)
Intermolecular bonding system Jeopardy! 139. It keeps ammonia NH3 together as a liquid what is 140. It keeps Br2 bromine a liquid, but iodine I2 a solid what is 141. It keeps phosphorus trichloride PCl3 together as a liquid what is 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? 142. Explain the difference
between bond polarity and molecular polarity. Who has the guts to stand and orate this answer? 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 doesnt exhibit radial symmetry is A polar molecule water A non polar molecule CCl 4 polar. 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 to Queens
143. The bonds in ozone O3 resonate back and forth, they are exceptional bonds called RESONATING BONDS. In reality the bonds are both 1 sized! 144. Once and for all, with a Lewis dot diagram, and one sentence, explain how carbon monoxide bonds together. C O C O Its 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.
~ C O 145. 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 145. 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. O2 molecules have double polar covalent bonds No, double nonpolar 4. N2 molecules have double nonpolar covalent bonds No, triple nonpolar 5. H atoms can make single or double covalent bonds No, only single 6. O atoms must make double bonds ONLY No, in water they make 2 singles 7. H2O is sometimes a straight line molecule by shape No, always bent! 8. Molecules with polar bonds can never be nonpolar molecules No, CO2 , 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