today is tuesday, march 24 th, 2015 pre-class: what kind of bond joins the two oxygen atoms in...

Today is Tuesday,March 24th, 2015

Pre-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).

Stuff You Need:See Below

In This Lesson:Covalent

Bonding, VSEPR Theory, and

Intermolecular Forces

(Lesson 4 of 4)

Today’s Agenda

• Covalent/Molecular Bonding• Bond/Molecular Polarity• VSEPR Theory• Intermolecular Forces

• Where is this in my book?– P. 213 and following…

By 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.

What you need to know…

• …is on this checklist.

Reminder

HH

O

This is an atom This is a

molecule

Video Introduction

• Find your Bonding Overview sheets and have your video worksheets ready (on the Covalent Bonding side).

• It’s another Australian video!

Covalent 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.

Characteristics 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.

Ionic 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.

The Official Way

• Remember 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!

Determining Bond Type

• Use the file Periodic Table – Electronegativity.• Compare the electronegativity values of the

two atoms.• Ionic: Difference ≥ 1.67– Sometimes people use numbers as high as 2.0.

• Polar Covalent: 1.66 ≥ Difference ≥ 0.4– This means that one atom (the more

electronegative one) “hogs” the shared electrons.• Non-Polar Covalent: 0.3 ≥ Difference– The atoms play nice.

The Unofficial Way

• Because 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.

Identify the Bond Types• NaCl

– Ionic (electronegativity difference of 2.1)

• F2

– Non-Polar Covalent (electronegativity difference of 0)

• MgCl2

– Ionic (electronegativity difference of 1.8)• BaO

– Ionic (electronegativity difference of 2.6)

• CH4

– Polar Covalent (electronegativity difference of 0.4)

• O2

– Non-Polar Covalent (electronegativity difference of 0)

Aside: 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.

The Octet Rule

• In 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.

Diatomic Molecules

• Recall that seven elements on the periodic table naturally bond to themselves.

• Elements that exist diatomically:– Br-I-N-Cl-H-O-F– Br2, I2, N2, Cl2, H2, O2, F2

• These elements are all bonded covalently to themselves.

The 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.

Facebook?

• 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.

http://infosthetics.com/archives/facebook_graph2.jpg

Facebook

Central Atoms

• In 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.

Now, 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.

Methods 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.

Single Bonds

• Single bonds involve two shared electrons (or one pair).– Single bonds are sometimes drawn with a single line.

F

F

1s

1s 2s

2s 2p

2p

Each has seven valence

electrons.

F F

Single Bond Example

One pair of electrons is

shared.FF

Think of it another way…

FF Shared

Electrons

Lewis Structure

Think of it another way…

F FShared

Electrons

Structural Formula

Covalent Bonding Practice

• Try this example: Hydrochloric Acid– HCl

H ClClH

Covalent Bonding Practice

• Try this example: Iodine– I2

I III

Covalent Bonding Practice

• Try this example: Dihydrogen Monoxide– H2O

HO HO

H H

Notice 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…

Covalent Bonding Practice

• Now it’s your turn.• Covalent Bonding Worksheet (first page) #2, 4,

5, 7, 9, 10

H

Unshared Pairs

• 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.

O H

There are two unshared pairs

of electrons.

Unshared Pairs

• Ammonia (NH3) has one unshared pair of electrons.– The electrons are on the nitrogen atom.

Unshared Pairs Practice

• How many unshared pairs are on the methane (CH4) molecule?– None.

Unshared Pairs Practice

• How many unshared pairs are on the carbon tetrachloride molecule?– None (carbon is the

central atom).

ClC ClCl

Cl

Double Bonds

• Double bonds involve four shared electrons (or two pairs).– Double bonds are sometimes drawn with double lines.

O

O

1s

1s 2s

2s 2p

2p

Each has six valence

electrons.

O O

Double Bond Example

Two pairs of electrons are

shared.OO

Covalent Bonding Practice

• Try this one: Carbon Dioxide– CO2

C OO

OCO

Covalent Bonding Practice

• Now it’s your turn.• Covalent Bonding Worksheet (first page) #3,

12

Triple Bonds

• Triple bonds involve six shared electrons (or three pairs).– Triple bonds are sometimes drawn with triple lines.

N

N

1s

1s 2s

2s 2p

2p

Each has five valence

electrons.

N N

Triple Bond Example

Three pairs of electrons are

shared.NN

Covalent Bonding Practice

• Now it’s your turn.• Covalent Bonding Worksheet #1• Covalent Bonding Worksheet (Reverse)– WARNING: I have done something a little sneaky

on Page 2.

Drawing 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.

Covalent Bonding Tips and TricksJot it like you mean it!

1. Hydrogen and halogens will only ever make one bond. Connect them with single bonds first.

2. Make sure you don’t use too many electrons! You can only use as many as the atoms bring.

3. Follow the HONC Rule when possible.4. Single bonds count as 2 electrons, double

count as 4, triple count as 6.5. DO NOT MAKE A QUADRUPLE BOND EVER!

CH3Cl: Method #1

HC HH

Cl

• Make carbon the central atom (it wants four bonds).

• Add up the available valence electrons:– C = 4, H = 1 (3), Cl = 7– TOTAL = 14.

• Join peripheral atoms to the central atoms with electron pairs.

• Complete octets on atoms other than hydrogen with remaining electrons.

• Make carbon the central atom (it wants four bonds).

• Add up the total electrons needed:– C = 8, H = 2 (3), Cl = 8– TOTAL = 22.

• Subtract available electrons:– C = 4, H = 1 (3), Cl = 7– TOTAL = 14.– 22 - 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.

CH3Cl: Method #2

HC HH

Cl

N

Practice Molecule: HCN

CHNCH

Practice

• Covalent Bonding and VSEPR Theory Worksheet– #1-3

• Covalent Bonding and VSEPR Theory Worksheet– #4-5

VSEPR Theory

• So 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?

VSEPR Theory

• VSEPR 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.

VSEPR 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.

VSEPR Theory

• To use the VSEPR Chart I gave you, do this:1. Draw the dot structure.2. Find:– The central atom.– The number of atoms bonded to the central atom.– The number of lone pairs on the central atom.

3. Then use the VSEPR Chart to find its shape.

VSEPR Chart Example: H2O

• We know H2O looks like this:

1. According to the HONC rule, O forms two bonds and H only forms one, so O is the central atom.

2. Oxygen has two bonded atoms and two unshared pairs of electrons.

3. The chart tells us that this is molecule takes on the “bent” geometry.

HO H

VSEPR Geometry: Linear

• 1 Bonded Atom or 2 Bonded Atoms with 0 Lone Pairs.– Example: CO2

C OO

VSEPR Geometry: Trigonal Planar

• 3 Bonded Atoms, 0 Lone Pairs.– Example: BBr3

BBr Br

Br

Wait, what?

• 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.

AlBr Br

Br

Covalent 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.

– ;)

BF F

F

VSEPR Geometry: Trigonal Pyramidal

• 3 Bonded Atoms, 1 Lone Pair.– Example: PH3

P

H HH

VSEPR Geometry: Bent

• 2 Bonded Atoms, 1 or 2 Lone Pairs.– Example: H2O

O

H H

VSEPR Geometry: Tetrahedral

• 4 Bonded Atoms, 0 Lone Pairs.– Example: CH4

C

H HH

H

Other Crazy Shapes

Seesaw

T-Shaped

Square Planar

Octahedral

Square Pyramidal

Trigonal Bipyramidal

VSEPR Summary

Bonded Atoms Lone Pairs Shape

1 0, 1, 2, or 3 Linear

2 0 Linear

2 1 or 2 Bent

3 0 Trigonal Planar

3 1 Trigonal Pyramidal

4 0 Tetrahedral

Practice

• Find 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.– #6

• Now add the geometry to #4 and 5 from Covalent Bonding and VSEPR Theory.

Practice

• Now try:– Covalent Bonding and VSEPR Theory Worksheet• Reverse – Bottom section

Practice

• Lab – Molecules I

Aside: Empty Space

• Remember 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.

Review

• Quick – what’s the shape?

Practice

• Now try:– Covalent Bonding and VSEPR Theory Worksheet• Reverse – Top section

Lastly…

• Finally, let’s do Lewis Structures for covalent ions.– Like polyatomic ions from waaaay back when…

Polyatomic Ion Lewis StructuresOr other really tricky ones…

• Procedure:1. Determine total valence electrons.2. Draw single bonds between all atoms.3. Give bonded atoms octets (except H).4. 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.

Electrons Left: 24Electrons Left: 0Electrons Left: 34

• 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…

• 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.

2-

OC

OC

O

OExample: Oxalate (C2O4

2-)

Then put it all in brackets with a charge indicated.

BIG 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.

Practice

• Covalent Bonding and VSEPR Theory Worksheet (Reverse, Middle Section)– Letters D, E, F, and G are challenging. Leave F for

last!

Covalent Bonding

• Remember that there are two types of covalent bonding:

• Polar Covalent– Electrons are not shared equally.– Electrons are pulled closer to one atom than

another because one atom is more electronegative.• Non-Polar Covalent– Electrons are shared equally.– Pulls from both atoms are roughly equal strength.

Water Has Polar Bonds• Let’s think of it this way. You tell me which of these

two guys is gonna hold onto his electrons more… (draw them too)

Meet Hydrogen – Poor guy’s only 2.1 Paulings in

electronegativity.

OMG OXYGEN!3.5 PAULINGS

ELECTRONEGATIVITY!!!1

But…since oxygen is such a strong atom, it doesn’t share electrons nicely.Having all that negative charge built up around oxygen causes that end of the molecule to take on a negative charge…

…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…

…which causes neighboring water molecules to form hydrogen bonds with the first one.This is the source of cohesion/surface tension, and why it hurts to do a belly-flop.

In a diagram…

HHO-

Meet water: (say hi)There are 8 total valence electrons in the molecule.

- - -----

-+

+

HH

O-++

HH

O

-+ +Hydrogen

Bond

Covalent Bond

Polar Covalent Bonds

• As 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: +

Polar Covalent Bonds

• Because 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.

http://www.chem.csustan.edu/chem2000/Exp1/water.gif

http://images.wikia.com/psychology/images/9/96/Water_drops_on_spider_web.jpghttp://www.factzoo.com/sites/all/img/reptiles/jesus-lizard-running-on-water-basilisk.jpg

Non-Polar Covalent Bonds

• Electrons shared (fought-over?) equally.• Example: H2 (or an H-H bond)

Bonds 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…

A figurative look at it?

• Think of when a child pulls on his/her parent’s arm to lead him/her somewhere:

ParentKid

This is a polar bond,

because “Parent” is

likely stronger than

“Kid.”

Other Kid

This is a polar molecule, because

the pulls do not cancel each other.

Add another kid with

equal strength and

you have another polar

bond.

This is a non-polar molecule, because the pulls do cancel

each other.

But if “Other Kid” is

ripped, the pulls may not cancel out.

This is a polar molecule, because

the pulls do not cancel each other.

Polar Bonds, Non-Polar Molecules

ClC ClCl

Cl +-

-

-

-

• 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.

Identifying 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: PH3

– Electronegativity 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*.

P

H HH

*Up for debate.

H2O – Polar or Non-Polar?

http://www.middleschoolchemistry.com/img/content/multimedia/chapter_5/lesson_1/water_molecule.jpg

• 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.

+

-

Another Look [sketch me]

• Let’s look at the “pulling” going on in a water molecule again:

HH

O3.5

2.1 2.1

O pulls downward on

electrons more strongly than H pulls upward…

…making O slightly negative……and H slightly

positive.KEY: To show this, we add a “plus” to

the arrow.-

+ +

Polar or Non-Polar?

• Carbon dioxide (CO2)– Non-polar; linear molecule

• Ammonia (NH3)– Polar; trigonal pyramidal molecule

• Chlorine (Cl2)– Non-polar; linear molecule

Practice

• Lab – Molecules I– Front: #2, 4– Back: “Molecule Polarity” Column

Aside: 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.

http://i1.wp.com/samhutchings.co/wp-content/uploads/2013/04/Radarange_first.jpg?resize=680%2C500

Intermolecular Forces

• Water in Space!

Intermolecular Forces

• Time for another Australian video!• Find your video worksheets and turn to the

reverse side for Intermolecular Forces.

Intermolecular Forces

• Intermolecular forces are interactive forces between molecules. Here are three, in order of increasing strength:– WEAKEST: London Dispersion Forces

• Occur: Always• Part of Van der Waals Forces

– Dipole-Dipole Forces• Occur: With two polar molecules.• Part of Van der Waals Forces

– STRONGEST: Hydrogen Bonding Forces• Occur: When H is bonded to O, N, or F on one molecule and

there’s an O, N, or F on another molecule.

London Dispersion Forces Demo

• I need a nucleus or two!

London 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.

London Dispersion Forces

• Part of Van der Waals Forces:

Dipole-Dipole Forces

• Also 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!

Van der Waals Forces

• Where do we see van der Waals forces in nature?– Geckos, for one!

http://duende.uoregon.edu/~hsu/blogfiles/gecko.jpg http://cdn.physorg.com/newman/gfx/news/2005/gecko.jpg

Geckos

Hydrogen 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.

Hydrogen 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.

http://academic.brooklyn.cuny.edu/biology/bio4fv/page/image12.gif

Intermolecular 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 BF3

– H2O Bent, BF3 Trigonal Planar

– H2O Polar, BF3 Non-Polar– London Dispersion, H-Bonds

• HCN and CH3F– HCN Linear, CH3F Tetrahedral

– HCN Polar, CH3F Polar– London Dispersion, Dipole-

Dipole

Important Note

• Generally 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.

Intermolecular Forces Practice

• Intermolecular Forces Worksheet

Bond 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.

Bond Dissociation Energies (kJ/mol)• H—H

• 436• H—C

• 393• H—N

• 391• H—O

• 464• H—F

• 568• H—Cl

• 432

• H—Br• 366

• H—I• 298

• C—C• 347

• C—O• 360

• C—N• 308

• C—F• 488

• C—Cl• 330

• C—S• 272

• C—Br• 288

• C—I• 216

• N—N• 170

• O—O• 145

• F—F• 158

• Cl—Cl• 243

• Br—Br• 192

• I—I• 151

• C=C• 657

• C=O• 736

• O=O• 498

• C≡C• 839

• N≡N• 945

Bond Dissociation Examples

• CH4

• 393 * 4 = 1572 kJ/mol

• CO2

• 736 * 2 = 1472 kJ/mol

• H2O• 464 * 2 = 928 kJ/mol

• Br2

• 192 kJ/mol

• I2

• 302 kJ/mol

• O2

• 498 kJ/mol• O—O• 145 kJ/mol

• NH3

• 391 * 3 = 1173 kJ/mol

Polystyrene Dissolution Demo

• Styrofoam (non-polar)• Acetone (non-polar)• Permanent marker (to draw the wicked witch

or something)

Bond Dissociation Energies

• Styrofoam (called Polystyrene)• Chemical Formula: C8H8

• Calculate 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/molTotal

6767 kJ/mol

Resonance Example

• Let’s draw the nitrate ion:– NO3

-

• 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?

N

O O

O

Resonance in Nitrate – NO3-

N

O O

O

N

O O

O

N

O O

O

Resonance

• Resonance 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.

N

O O

O

Resonance in Benzene (C6H6)

Aside: 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:

S

O O

O Double

Bond

Coordinate Bond

Coordinate Bond

Aside: 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.

Resonance in Nitrate – NO3-

N

O O

O

N

O O

O

N

O O

O

What’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.

And the answer to #6?With coordinate bonds: As a single resonance

structure:

S

O O

O

S

O O

O

Closure

• Wow.• We’ve done a lot of stuff throughout this

lesson.• Questions?

Okay, 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…

Orbital Hybridization

Groups Attached to

Central AtomGeometry Hybridization

2 Linear sp3 Trigonal Planar sp2

4 Tetrahedral sp3

5 Trigonal Bipyramidal sp3d

6 Octahedral sp3d2

Hybrid Bonds

Bond Type Hybrid Bond TypesSingle 1 Sigma Bond*

Double 1 Sigma Bond*1 Pi Bond**

Triple 1 Sigma Bond*2 Pi Bonds**

*Overlapping of two sp hybrid orbitals.**Overlapping of two non-hybridized p orbitals.