Presentation on theme: "Today is Tuesday, March 24th, 2015"— Presentation transcript:
1Today is Tuesday, March 24th, 2015 In This Lesson:Covalent Bonding, VSEPR Theory, and Intermolecular Forces(Lesson 4 of 4)Stuff You Need:See BelowToday is Tuesday, March 24th, 2015Pre-Class:What kind of bond joins the two oxygen atoms in oxygen gas – O2?What kind of bond joins the carbon and oxygen atoms in carbon dioxide – CO2?Please take a video worksheet from the Turn-In Box and get a paper towel. Also find your bonding overview worksheets (the big grid one).
2Today’s Agenda Covalent/Molecular Bonding Bond/Molecular Polarity VSEPR TheoryIntermolecular ForcesWhere is this in my book?P. 213 and following…
3By the end of this lesson… You should be able to describe the nature of a molecular bond.You should be able to describe the shape of a molecule based on its constituent atoms.You should be able to identify forces between molecules based on the nature of the participating molecules.You should be able to draw resonance structures of molecules.
6Video IntroductionFind your Bonding Overview sheets and have your video worksheets ready (on the Covalent Bonding side).It’s another Australian video!
7Covalent Bonds: The Important Stuff Covalent bonds are sometimes called molecular bonds.Covalent bonds generally occur between nonmetals and nonmetals (or metalloids).Electron pairs are shared (or fought-over).Bonds can be polar or non-polar.The term “molecule” is used only for covalent compounds.Ionic compounds sometimes get the term formula unit.
8Characteristics of Molecular Compounds Poor conductors of electricity.Low melting points.Polar compounds can dissolve in water.Non-polar compounds cannot dissolve in water.Many are gases at room temperature.
9Ionic or Covalent? But wait, you say. How can you determine whether a compound is ionic or covalent?There are two ways.The “official” way:By electronegativity.The “unofficial” way:By location on the periodic table.
10The Official WayRemember that electronegativity is the ability of an atom to attract electrons in a covalent compound.Why are noble gases not on here?They don’t form compounds!
11Determining Bond Type Use the file Periodic Table – Electronegativity. Compare the electronegativity values of the two atoms.Ionic: Difference ≥ 1.67Sometimes people use numbers as high as 2.0.Polar Covalent: 1.66 ≥ Difference ≥ 0.4This means that one atom (the more electronegative one) “hogs” the shared electrons.Non-Polar Covalent: 0.3 ≥ DifferenceThe atoms play nice.
12The Unofficial WayBecause ionic compounds occur between cations and anions, they often occur between elements on opposite sides of the periodic table.Consider both ways when making the distinction between bonds.
13Identify the Bond Types NaClIonic (electronegativity difference of 2.1)F2Non-Polar Covalent (electronegativity difference of 0)MgCl2Ionic (electronegativity difference of 1.8)BaOIonic (electronegativity difference of 2.6)CH4Polar Covalent (electronegativity difference of 0.4)O2
14Aside: Bond Type Exceptions? You will not need to know them for this class, but keep in mind that there are exceptions to this bond type rule.A notable one is compounds involving Beryllium (Be) – they tend to be covalent despite Beryllium not achieving an octet.BeCl2, for example, is a covalent compound as confirmed by experimental evidence involving boiling point and behavior.However, it still likes to form things that look like crystal lattices when it’s a solid. Weird.
15The Octet RuleIn ionic compounds, electrons are transferred such that atoms achieve the octet, either by losing or gaining.In covalent compounds, electrons are shared (or fought over) so that each atom has eight.The key?Whichever type of compound it is, everyone either needs to lose all their valence electrons or share/gain to eight.
16Diatomic MoleculesRecall that seven elements on the periodic table naturally bond to themselves.Elements that exist diatomically:Br-I-N-Cl-H-O-FBr2, I2, N2, Cl2, H2, O2, F2These elements are all bonded covalently to themselves.
17The HONC Rule* *Can be broken on occasion. The HONC Rule is an easy way to show how many bonds form for an element.Hydrogen and Halogens form one covalent bond.Oxygen (and sulfur) form two covalent bonds.Either two single bonds or one double bond.Nitrogen (and phosphorus) form three covalent bonds.Either a triple bond, a double and single bond, or three single bonds.Carbon (and silicon) form four covalent bonds.Either a triple and a single bond, two double bonds, a double and two single bonds, or four single bonds.
18Facebook?Back in the day, Facebook had a “visualizer” for your friends.Basically, it made a web of friend connections.Those with more connections tended to be in the inner part of the web, those with fewer were on the outside.
20Central AtomsIn a molecule, the central atom is the one that forms the most bonds.According to the HONC Rule, Carbon and Silicon (and others in their group) are most likely to be the central atom.Similarly, hydrogens and halogens are rarely the central atoms.
21Now, onto the bonding…As you know, Lewis Structures (dot notation? Remember?) show the arrangement of valence electrons in an atom.In covalent bonds, shared/fought over electrons can be shown by two dots ( : ) or a single line ( - ).IMPORTANT: Each single line (-) connects two dots, so you should count them as two electrons.
22Methods of Visualization If you’re using strictly dots, you’re drawing the Lewis Dot Structure.If you’re using lines, you’re drawing the Structural Formula.Throughout this lesson, you’ll learn many ways to accomplish one goal. Use what works for you and/or the problem.
23Single Bonds Single bonds involve two shared electrons (or one pair). Single bonds are sometimes drawn with a single line.
24F F F F Single Bond Example F F Each has seven valence electrons. One pair of electrons is shared.FF
25Think of it another way… SharedElectronsLewis Structure
26Think of it another way… SharedElectronsFStructural Formula
27Covalent Bonding Practice Try this example: Hydrochloric AcidHClHClHCl
28Covalent Bonding Practice Try this example: IodineI2IIII
29Covalent Bonding Practice Try this example: Dihydrogen MonoxideH2OOHOHHH
30Notice something…In each of the examples we did, I drew in the non-bonding pairs of electrons.As we saw in the video, these electrons have a strong influence on the shape of the molecule even without bonding to another atom.More to come…
31Covalent Bonding Practice Now it’s your turn.Covalent Bonding Worksheet (first page) #2, 4, 5, 7, 9, 10
32There are two unshared pairs of electrons. In our all examples (but let’s return to H2O), there are pairs of electrons that are not shared.There are known as unshared pairs, lone pairs, or non-bonding pairs.IMPORTANT: Unshared pairs only count on the central atom (the one with the most bonds).You should still draw them, though.OHHThere are two unshared pairs of electrons.
33Unshared Pairs Ammonia (NH3) has one unshared pair of electrons. The electrons are on the nitrogen atom.
34Unshared Pairs Practice How many unshared pairs are on the methane (CH4) molecule?None.
35Unshared Pairs Practice How many unshared pairs are on the carbon tetrachloride molecule?None (carbon is the central atom).ClClCClCl
36Double BondsDouble bonds involve four shared electrons (or two pairs).Double bonds are sometimes drawn with double lines.
37O O O O Double Bond Example O O Each has six valence electrons. Two pairs of electrons are shared.OO
38Covalent Bonding Practice Try this one: Carbon DioxideCO2OCOOCO
39Covalent Bonding Practice Now it’s your turn.Covalent Bonding Worksheet (first page) #3, 12
40Triple BondsTriple bonds involve six shared electrons (or three pairs).Triple bonds are sometimes drawn with triple lines.
41N N N N Triple Bond Example N N Each has five valence electrons. Three pairs of electrons are shared.NN
42Covalent Bonding Practice Now it’s your turn.Covalent Bonding Worksheet #1Covalent Bonding Worksheet (Reverse)WARNING: I have done something a little sneaky on Page 2.
43Drawing Lewis Structures Bigger molecules can make for complicated Lewis Structure drawings.Imagine drawing the Lewis Structure for CH3Cl.Thankfully, there are two systematic ways to do it.
44Covalent Bonding Tips and Tricks Jot it like you mean it! Hydrogen and halogens will only ever make one bond. Connect them with single bonds first.Make sure you don’t use too many electrons! You can only use as many as the atoms bring.Follow the HONC Rule when possible.Single bonds count as 2 electrons, double count as 4, triple count as 6.DO NOT MAKE A QUADRUPLE BOND EVER!
45CH3Cl: Method #1Make carbon the central atom (it wants four bonds).Add up the available valence electrons:C = 4, H = 1 (3), Cl = 7TOTAL = 14.Join peripheral atoms to the central atoms with electron pairs.Complete octets on atoms other than hydrogen with remaining electrons.HClCHH
46CH3Cl: Method #2Make carbon the central atom (it wants four bonds).Add up the total electrons needed:C = 8, H = 2 (3), Cl = 8TOTAL = 22.Subtract available electrons:C = 4, H = 1 (3), Cl = 7TOTAL = 14.= 8.Divide by two. (8/2 = 4)Draw in 4 total bonds between atoms, then complete octets.Remember, double bonds count as two and triple as three.HClCHH
48PracticeCovalent Bonding and VSEPR Theory Worksheet#1-3#4-5
49VSEPR TheorySo dot structures are just peachy, but they leave out something important.Something like a third dimension.The thing is, thinking about things in three dimensions can be difficult.So…let’s introduce the next topic with a video:TED: George Zaidan and Charles Morton – What is the Shape of a Molecule?
50VSEPR TheoryVSEPR Theory (pronounced “vesper”) stands for Valence Shell Electron Pair Repulsion Theory.You can think of it as a theory that helps scientists predict the shape of a molecule.Because electrons involved in bonding pairs or lone pairs are both negatively charged, they repel one another and get as far apart as possible.
51VSEPR Theory Chart Treat this VSEPR Theory Chart like gold. If you need another, it’s on my website.This is invaluable and is basically your lifesaver.For reals.It summarizes a lot of information that can’t be put into a PowerPoint and I’ll be referring to it frequently.Unfortunately, it will not be available to you on the test.
52VSEPR Theory To use the VSEPR Chart I gave you, do this: Draw the dot structure.Find:The central atom.The number of atoms bonded to the central atom.The number of lone pairs on the central atom.Then use the VSEPR Chart to find its shape.
53VSEPR Chart Example: H2O We know H2O looks like this:According to the HONC rule, O forms two bonds and H only forms one, so O is the central atom.Oxygen has two bonded atoms and two unshared pairs of electrons.The chart tells us that this is molecule takes on the “bent” geometry.OHH
54VSEPR Geometry: Linear 1 Bonded Atom or 2 Bonded Atoms with 0 Lone Pairs.Example: CO2OCO
56Wait, what? Al Br Try drawing the dot structure for AlBr3. Notice anything?As you probably could have guessed, there are exceptions to the octet rule occasionally.The most common exceptions are compounds that have only single bonds formed with Group 3A cations.These compounds do not obey the octet rule and take on trigonal planar geometry.
57Covalent Bonding Worksheet Take a look at #9 on the reverse of your Covalent Bonding Worksheet.If this were Boron (and thus not an ionic compound), this would be a molecular compound and would exhibit the same geometry.The reason it’s ionic is because of the electronegativity difference between Al and F.;)FBFF
62VSEPR Summary Bonded Atoms Lone Pairs Shape 1 0, 1, 2, or 3 Linear 2 1 or 2Bent3Trigonal PlanarTrigonal Pyramidal4Tetrahedral
63PracticeFind your Covalent Bonding Worksheet and write in the geometry of each covalently-bonded molecule.Example: #1 (N2) is linear because it has one bonded atom and one lone pair.Now try Covalent Bonding and VSEPR Theory Worksheet.#6Now add the geometry to #4 and 5 from Covalent Bonding and VSEPR Theory.
64Practice Now try: Covalent Bonding and VSEPR Theory Worksheet Reverse – Bottom section
66Aside: Empty SpaceRemember earlier in the year when I told you that we never really touch anything?As in, repulsive forces between atoms keep us from actually coming in contact with them so we just get really really close?Well, this VSEPR stuff, the repulsion of electrons, is what’s causing that.VSEPR Theory is part of the reason you can’t walk through a wall right now.You actually have plenty of space to sneak your atoms’ nuclei through the wall’s – if it weren’t for the repulsive forces.
68Practice Now try: Covalent Bonding and VSEPR Theory Worksheet Reverse – Top section
69Lastly… Finally, let’s do Lewis Structures for covalent ions. Like polyatomic ions from waaaay back when…
70Polyatomic Ion Lewis Structures Or other really tricky ones… Procedure:Determine total valence electrons.Draw single bonds between all atoms.Give bonded atoms octets (except H).Satisfy the octet rule for the central atom(s).Any leftover electrons go to the central atom.If the central atom has six electrons, it double bonds to a bonded atom with a lone pair.If the central atom has four electrons, it triple bonds to a bonded atom with a lone pair, or makes two double bonds.BIG HUGE IMPORTANT NOTE!The HONC rule can be broken, and will always be for ions but never for hydrogen or halogens.
71Example: Oxalate (C2O42-) Electrons Left: 24Electrons Left: 34Electrons Left: 0Example: Oxalate (C2O42-)2-Draw a single bond between all atoms.Give bonded atoms octets.If a central atom has only six total electrons, it double bonds to a bonded atom with a lone pair.Find the total e-:Each carbon has 4 and each oxygen has 6, so 32 total.The overall charge is 2-, so there are 2 more electrons.Grand total 34.Start filling in…OOCCOOThen put it all in brackets with a charge indicated.
72BIG IMPORTANT NOTE DEUX The method just described works well for challenging dot structures.In case of emergency, use this technique, but don’t forget to count your electrons.
73PracticeCovalent Bonding and VSEPR Theory Worksheet (Reverse, Middle Section)Letters D, E, F, and G are challenging. Leave F for last!
74Covalent BondingRemember that there are two types of covalent bonding:Polar CovalentElectrons are not shared equally.Electrons are pulled closer to one atom than another because one atom is more electronegative.Non-Polar CovalentElectrons are shared equally.Pulls from both atoms are roughly equal strength.
75Water Has Polar BondsLet’s think of it this way. You tell me which of these two guys is gonna hold onto his electrons more… (draw them too)OMG OXYGEN!3.5 PAULINGS ELECTRONEGATIVITY!!!1Meet Hydrogen – Poor guy’s only 2.1 Paulings in electronegativity.
76- - - + + + + O O O H H H In a diagram… - - - - - - - - There are 8 total valence electrons in the molecule.Meet water: (say hi)This is the source of cohesion/surface tension, and why it hurts to do a belly-flop.…which causes neighboring water molecules to form hydrogen bonds with the first one.…and all that’s left on the other end are the hydrogen protons, causing that end of the molecule to take on a positive charge…Having all that negative charge built up around oxygen causes that end of the molecule to take on a negative charge…But…since oxygen is such a strong atom, it doesn’t share electrons nicely.+HO-Covalent Bond+HO-++HOHydrogen Bond---------
77Polar Covalent BondsAs a result of that big oxygen dude hoggin’ all the negative electrons (and being more electronegative), his end of the water molecule takes on a slightly negative charge.We use the (lower case) delta symbol to indicate a negative charge: -Because poor ol’ Hydrogen’s doesn’t see much of the electron action (and being less electronegative), his end gets a slightly positive charge (due to the hydrogen protons).We use the (lower case) delta symbol to indicate a positive charge: +
78Polar Covalent BondsBecause water molecules have poles, the positive end of one water molecule (hydrogen) can stick to the negative end of another (oxygen).This is what allows water to undergo adhesion and cohesion.
80Bonds and Molecules Now we separate the molecules from the bonds. As we have learned, individual bonds can be polar.However, as in the lab, sometimes molecules made only of polar bonds can be non-polar overall.Here’s what I mean…
81This is a polar molecule, because the pulls do not cancel each other. This is a non-polar molecule, because the pulls do cancel each other.A figurative look at it?Think of when a child pulls on his/her parent’s arm to lead him/her somewhere:This is a polar bond, because “Parent” is likely stronger than “Kid.”But if “Other Kid” is ripped, the pulls may not cancel out.Add another kid with equal strength and you have another polar bond.KidParentOther Kid
82Polar Bonds, Non-Polar Molecules Carbon tetrachloride (CCl4) has 4 polar bonds.Electronegativity difference of 0.5 in each direction.However, because all chlorine atoms are pulling equally on the electrons, the overall molecule remains symmetrical, and thus non-polar.Cl-ClCCl+--Cl-
83Identifying Polar Molecules If the “pulls” in a molecule are symmetrical and cancel each other, it’s non-polar.Even if the bonds are polar.If they are asymmetrical, it is polar.Example: PH3Electronegativity difference: 2.1 (P) – 2.1 (H) = 0.So all bonds are non-polar.BUT! There is an unshared pair on P, making the molecule assymetrical (trigonal pyramidal geometry), leading to a polar molecule*.PHHH*Up for debate.
84H2O – Polar or Non-Polar? - + Here’s another model of water. We know water’s polar, but it appears to be symmetrical, right?However, when you look at the direction of the pulls, it is clearly not pulling equally in all directions.So, to determine molecular polarity, draw the “pulls” from polar bonds and see if they cancel out.-+
85Another Look [sketch me] Let’s look at the “pulling” going on in a water molecule again:2.1++2.1HO…making O slightly negative…O pulls downward on electrons more strongly than H pulls upward……and H slightly positive.KEY: To show this, we add a “plus” to the arrow.3.5-
86Polar or Non-Polar? Carbon dioxide (CO2) Ammonia (NH3) Chlorine (Cl2) Non-polar; linear moleculeAmmonia (NH3)Polar; trigonal pyramidal moleculeChlorine (Cl2)
88Aside: Microwave Ovens Microwaves work by heating water molecules in food.Since the plate usually doesn’t have water, it doesn’t get warm.How do they heat the water? With polarity!The oven induces an electromagnetic field that causes the polar water molecules to orient in a certain direction.Then the microwave changes the orientation of the field and the molecules need to change their orientation too.The shifting back and forth and resulting friction creates heat and deliciousness.
90Intermolecular Forces Time for another Australian video!Find your video worksheets and turn to the reverse side for Intermolecular Forces.
91Intermolecular Forces Intermolecular forces are interactive forces between molecules. Here are three, in order of increasing strength:WEAKEST: London Dispersion ForcesOccur: AlwaysPart of Van der Waals ForcesDipole-Dipole ForcesOccur: With two polar molecules.STRONGEST: Hydrogen Bonding ForcesOccur: When H is bonded to O, N, or F on one molecule and there’s an O, N, or F on another molecule.
92London Dispersion Forces Demo I need a nucleus or two!
93London Dispersion Forces The weakest type of intermolecular force.It results from the motion of electrons around the nucleus.When, momentarily, electrons in one atom happen to be on the same side of the atom, a transient dipole forms.It’s like a temporary polar molecule.As a result, the molecule has a polar end and non-polar end, and nearby molecules can also gain transient dipoles, causing them to attract each other.
94London Dispersion Forces Part of Van der Waals Forces:
95Dipole-Dipole ForcesAlso part of Van der Waals forces, dipole-dipole forces occur when two polar molecules come together.The positive end of one polar molecule is attracted to the negative end of another.Polar molecules are called dipoles, after all!
96Van der Waals Forces Where do we see van der Waals forces in nature? Geckos, for one!
98Hydrogen Bonding The strongest force between neutral molecules. Also known as “when Hydrogen cheats on its molecule.”When is there H-bonding? (NEED BOTH CONDITIONS)On one molecule, H is bonded to O, N, or F.Another molecule with O, N, or F is nearby.Mostly because H only has two electrons.
99Hydrogen Bonding Explained When is there H-bonding (long answer)?When H, small little atom that it is, bonds to a strongly electronegative atom like oxygen, nitrogen, or fluorine, none of the electrons hang out near it at all, making it positive.Similarly, the O, N, or F on another atom usually has all the electrons hanging out near it, making it negative.As a result, the H on one molecule and the O, N, or F on another are attracted to one another. Like a specific dipole-dipole interaction.Or some kind of extra-molecular couple.
100Intermolecular Forces Practice Draw the pair and identify:The shape of each.Whether each is polar or non-polar.The intermolecular forces between the two.H2O and BF3H2O Bent, BF3 Trigonal PlanarH2O Polar, BF3 Non-PolarLondon Dispersion, H-BondsHCN and CH3FHCN Linear, CH3F TetrahedralHCN Polar, CH3F PolarLondon Dispersion, Dipole-Dipole
101Important NoteGenerally speaking, the more intermolecular forces between two molecules, the higher the boiling and melting points.This is because the molecules hang onto each other more strongly.
102Intermolecular Forces Practice Intermolecular Forces Worksheet
103Bond Dissociation Energies For chemical reactions to occur, bonds must be broken.Bond dissociation energy is the amount of energy needed to break a covalent bond.Of one mole of the substance – more to come.To find the dissociation energy of the whole molecule, add the bond dissociation energies of each bond in the molecule.
106Polystyrene Dissolution Demo Styrofoam (non-polar)Acetone (non-polar)Permanent marker (to draw the wicked witch or something)
107Bond Dissociation Energies Styrofoam (called Polystyrene)Chemical Formula: C8H8Calculate the total bond dissociation energy for polystyrene.7 C-H Bonds:393 * 7 = 2751 kJ/mol4 C-C Bonds:347 * 4 = 1388 kJ/mol4 C=C Bonds:657 * 4 = 2628 kJ/molTotal6767 kJ/mol
108O N O O Resonance Example Let’s draw the nitrate ion: Add up total electrons available:24.Put single bonds everywhere.Complete octets on bonded atoms.Satisfy the octet rule for the central atom.Did you put your double bond somewhere else?ONOO
110ResonanceResonance occurs when more than one valid Lewis structure can be drawn for a particular molecule.The actual structure (or bond order) is an average of resonance structures.Additional bonds are shown as dotted lines, as in this image.ONOO
112Aside: Coordinate Covalent Bonds On your Covalent Bonding Worksheets, #6 on the front (SO3) is an example of a substance with a coordinate covalent bond.This is when there is a “donation” of an electron pair from one atom to another.Both electrons are donated from one atom.Here’s how to draw it:O Double BondSCoordinate Bond Coordinate BondOO
113Aside: Coordinate Covalent Bonds Note that coordinate covalent bonds (also known as dipolar bonds or a dative covalent bond) frequently occur in resonance structures.A good sign of this is when an atom has all its valence electrons shown as lone pairs, yet is still forming a covalent bond.Usually oxygen, it seems like, but could be others.Let’s redraw nitrate’s resonance structures to show coordinate covalent bonds.
115What’s the Big Idea?Coordinate covalent bonds do not function any differently from the kind of covalent bonding you already know how to draw.Drawing a coordinate bond simply allows you to specify the origin of the electrons that are being shared.You do not need to show coordinate bonds for this class, but it’s good to know about them for those of you with AP Chem in your futures.
116As a single resonance structure: With coordinate bonds: And the answer to #6?As a single resonance structure:With coordinate bonds:SOSO
117Closure Wow. We’ve done a lot of stuff throughout this lesson. Questions?
118Okay, one last thing…Should you decide to take AP Chemistry, you’ll also cover a little bit more of VSEPR Theory called hybridization.Hybridization occurs in covalent bonding, as electron repulsion changes the shape of orbitals (generally just the s and p ones, but sometimes d).When hybridization occurs, the orbitals re-shape to be more like a strange combination of both of them.When those orbitals form bonds with each other, the bond type is renamed as well.Here’s a short summary…
119Orbital Hybridization Groups Attached to Central AtomGeometryHybridization2Linearsp3Trigonal Planarsp24Tetrahedralsp35Trigonal Bipyramidalsp3d6Octahedralsp3d2
120Hybrid Bonds Bond Type Hybrid Bond Types Single 1 Sigma Bond* Double 1 Pi Bond**Triple2 Pi Bonds***Overlapping of two sp hybrid orbitals.**Overlapping of two non-hybridized p orbitals.