Chapter 1 A Review of General Chemistry Electrons Bonds and Molecular Properties David Klein Copyright 2015 John Wiley amp Sons Inc All rights reserved Klein Organic Chemistry 2e 11 Organic Chemistry ID: 642367
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Slide1
Organic Chemistry
Second Edition
Chapter 1A Review of General Chemistry: Electrons, Bonds, and Molecular Properties
David Klein
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e Slide2
1.1 Organic Chemistry
The study of carbon-containing molecules and their reactionsWhat happens to a molecule during a reaction?A collisionBonds break/formThe BIG question: WHY do reactions occur?We will need at least 2 semesters of your time to answer this questionFOCUS ON THE ELECTRONS
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e 1-2Slide3
1.1 Organic Chemistry
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e 1-3
Wöhler, 1928 Slide4
1.2 Structural Theory
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-4Slide5
1.2 Structural Theory
Atoms that are most commonly bonded to carbon include N, O, H, and halides (F, Cl, Br, I).With some exceptions, each element generally forms a specific number of bonds with other atomsPractice with SkillBuilder 1.1
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e 1-
5Slide6
A covalent bond is a PAIR of electrons shared between two atoms. For example…
1.3 Covalent Bonding
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Klein, Organic Chemistry 2e
1-6Slide7
How do potential energy and stability relate?
What forces keep the bond at the optimal length?1.3 Covalent Bonding
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Klein, Organic Chemistry 2e
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7Slide8
1.3 Atomic Structure
A review from General ChemistryProtons (+1) and neutrons (neutral) reside in the nucleusElectrons (-1) reside outside the nucleus.Some electrons are close to the nucleus and others are far away.
Look at carbon for example. Which electrons are the valence electrons? Why are valence electrons important?Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-8Slide9
You can always calculate the number of valence electron by analyzing the e- configuration. Or, for Group A elements only, just look at the Group number (Roman Numeral) on the periodic table
Practice with SkillBuilder 1.2
1.3 Counting Valence ElectronsCopyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-9Slide10
1.3 Simple Lewis Structures
For simple Lewis Structures…Draw the individual atoms using dots to represent the valence electrons.Put the atoms together so they share PAIRS of electrons to make complete octets. WHAT is an octet? Take NH
3, for example…Practice with SkillBuilder 1.3
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Klein, Organic Chemistry 2e
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1.3 Simple Lewis Structures
For simple Lewis Structures…Draw the individual atoms using dots to represent the valence electrons.Put the atoms together so they share PAIRS of electrons to make complete octets. WHAT is an octet? Try drawing the structure for C
2H2Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-11Slide12
1.4 Formal Charge
What term do we use to describe atoms with an unbalanced or FORMAL charge?How does formal charge affect the stability of an atom?Atoms in molecules (sharing electrons) can also have unbalanced charge, which must be analyzed, because it affects stabilityTo calculate FORMAL charge for an atom, compare the number of valence electrons that should be associated with the atom to the number of valence electrons that are actually associated with an atom
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e 1-12Slide13
1.4 Formal Charge
Consider the formal charge example below. Calculate the formal charge on each atom. orCarbon should have 4 valence electrons,
because it is in group IVA on the periodic table.Carbon actually has 8 valence electrons. It needs 8 for its octet, but only 4 count towards its charge. WHY?The 4 it actually has balance out the 4 it should have, so it does not have formal charge. Its neutral.
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Klein, Organic Chemistry 2e
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1.4 Formal Charge
Analyze the formal charge of the oxygen atom.
orOxygen should have 6 valence electrons, because it is in group VIA on the periodic table.It actually has 8 valence electrons. It needs 8 for its octet, but only 7 count towards its charge
. WHY?If it actually has 7, but it should only have 6, what is its formal charge?Practice with SkillBuilder
1.4Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e 1-
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1.5 Polar Covalent Bonds
Covalent bonds are electrons pairs that exist in an orbital shared between two atoms. What do you think that orbital looks like?Just like an atomic orbital, the electrons could be anywhere within that orbital region.What factors determine which atom in the bond will attract the shared electrons more?
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e 1-15Slide16
1.5 Polar Covalent Bonds
Covalent bonds are either polar or nonpolarNonpolar Covalent –bonded atoms share electrons evenlyPolar Covalent – One of the atoms attracts electrons more than the otherElectronegativity - how strongly an atom attracts shared electrons
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Klein, Organic Chemistry 2e
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1.5 Polar Covalent Bonds
Electrons tend to shift away from lower electronegativity atoms to higher electronegativity atoms. The greater the difference in electronegativity, the more polar the bond.
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Klein, Organic Chemistry 2e
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1.5 Polar Covalent Bonds
Can a bond have both covalent and ionic character?Practice with SkillBuilder 1.5Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e 1-18Slide19
1.6 Atomic Orbitals
General Chemistry reviewIn the 1920s, Quantum Mechanics was established as a theory to explain the wave properties of electronsThe solution to wave equations for electrons provides us with visual pictures called orbitals
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e 1-
19Slide20
1.6 Atomic Orbitals
General Chemistry reviewThe type or orbital be identified by its shapeAn orbital is a region where there is a calculated 90% probability of finding an electron. The remaining 10% probability tapers off as you move away from the nucleus
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-20Slide21
1.6 Atomic Orbitals
Electrons behave as both particles and waves. How can they be BOTH? Maybe the theory is not yet complete
The theory does match experimental data, and it has predictive capability.Like a wave on a lake, an electron’s wavefunction can be (+), (–), or ZERO. Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-21Slide22
1.6 Atomic Orbitals
Because they are generated mathematically from wavefunctions, orbital regions can also be (–), (+), or ZEROThe sign of the wave function has nothing to do with electrical charge.
In this p-orbital, there is a nodal plane. The sign of the wavefunction will be important when we look at orbital overlapping in bonds.Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-22Slide23
1.6 Atomic Orbitals
Electrons are most stable (lowest in energy) if they are in the 1s orbital?The 1s orbital is full once there are two electrons in it. Why
can’t it fit more?The 2s orbital is filled next. The 2s orbital has a node. WHERE?
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-23Slide24
Once the 2s is full, electrons fill into the
three degenerate 2p orbitalsWhere are the nodes in each of the 2p orbitals?
1.6 Atomic OrbitalsCopyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-24Slide25
1.6 Atomic Orbitals
Common elements and their electron configurationsPractice with SkillBuilder 1.6
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Klein, Organic Chemistry 2e 1-25Slide26
1.6 Atomic Orbitals
What are the rules that govern our placement of electrons ?Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-26Slide27
A bond occurs when atomic orbitals overlap
. Overlapping orbitals is like overlapping wavesOnly constructive interference results in a bond
1.7 Valence Bond Theory
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1-27Slide28
The bond for a H
2 molecule results from constructive interferenceWhere do the bonded electrons spend most of their time?
1.7 Valence Bond TheoryCopyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e
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1.8 Molecular Orbital Theory
Atomic orbital wavefunctions overlap to form MOs that extend over the entire molecule.MOs are a more complete analysis of bonds, because they include both constructive and destructive interference.The number of MOs created must be equal to the number of AOs that were used.
H2 MOsCopyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e 1-29Slide30
Why is the antibonding orbital higher in energy?
When the AOs overlap, why do the electrons go into the bonding MO rather than the antibonding MO?1.8 Molecular Orbital Theory
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e 1-
30Slide31
1.8 Molecular Orbital Theory
Imagine a He2 molecule. How would its MOs compare to those for H2?In general, if a molecule has all of
its bonding and antibonding MOs occupied, will it be stable or unstable?How would the energy of the He
2 compare to 2 He?Why does Helium exist in its atomic form rather than in molecular form?Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e
1-31Slide32
1.8 Molecular Orbital Theory
Consider TWO of the many MOs that exist for CH3BrThere are many areas of atomic orbital overlapNotice how the MOs extend over the entire moleculeEach picture below represents ONE orbital.
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Klein, Organic Chemistry 2e
1-32Slide33
1.8 Molecular Orbital Theory
How many electrons can fit into the areas represented?In the ground state, electrons occupy some MOs and not others, WHY?Depending on the circumstances, we will use both MO and valence bond theory to explain phenomena
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-33Slide34
1.9 Hybridized Atomic Orbitals
Given the electron configuration for C and H, imagine how their atomic orbitals might overlapWould such orbital overlap
yield methane?
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Klein, Organic Chemistry 2e
1-34Slide35
1.9 Hybridized Atomic Orbitals
To make methane, the C atom must have 4 atomic orbitals available for overlappingIf an electron is excited from the 2s to the 2p, will that make it suitable for making methane?
If four H atoms were to come in and overlap with the 2s and 2p orbitals, what geometry would the resulting methane have?Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e
1-35Slide36
1.9 Hybridized Atomic Orbitals
The carbon must undergo hybridization to form 4 equal atomic orbitalsThe atomic orbitals must be equal in energy to form four equal-energy symmetrical C-H bonds
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e 1-
36Slide37
1.9 Hybridized Atomic Orbitals
Should the shape of an sp3 orbital look more like an s or more like p orbital?
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Klein, Organic Chemistry 2e 1-37Slide38
1.9 Hybridized Atomic Orbitals
To make CH4, the 1s atomic orbitals of four H atoms will overlap with the four sp3 hybrid atomic orbitals of C
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Klein, Organic Chemistry 2e
1-38Slide39
1.9 Hybridized Atomic Orbitals
Draw a picture that shows the necessary atomic orbitals and their overlap to form ethane (C2H6).Draw a picture that shows the necessary atomic orbitals and their overlap to form water.Practice with conceptual checkpoint 1.19
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e 1-39Slide40
Consider ethene (ethylene).
Each carbon in ethene must bond to three other atoms, so only three hybridized atomic orbitals are needed
1.9 Hybridized Atomic Orbitals
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Klein, Organic Chemistry 2e
1-40Slide41
1.9 Hybridized Atomic Orbitals
An sp2 hybridized carbon will have three equal-energy sp2 orbitals and one unhybridized p orbital
Which is lower in energy, the sp2 or the p? Why?
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-41Slide42
1.9 Hybridized Atomic Orbitals
The sp2 atomic orbitals overlap to form sigma (σ) bonds
Sigma bonds provide maximum HEAD-ON overlapCopyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-42Slide43
1.9 Hybridized Atomic Orbitals
The unhybridized p orbitals in ethene form pi (π) bonds, SIDE-BY-SIDE overlap
Practice with conceptual checkpoint 1.20Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e
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1.9 Hybridized Atomic Orbitals
The unhybridized p orbitals in ethene form pi (π) bonds, SIDE-BY-SIDE overlap of p-orbitals giving both CONSTRUCTIVE and DESTRUCTIVE interference
MO theory shows the orbitals that result. Remember, red and blue regions are all part of the same orbital
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Klein, Organic Chemistry 2e 1-
44Slide45
1.9 Hybridized Atomic Orbitals
Why is sp2 hybridization not appropriate for methane (CH4)?
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e 1-45Slide46
1.9 Hybridized Atomic Orbitals
Consider ethyne (acetylene). Each carbon in ethyne must bond to two other atoms, so only two hybridized atomic orbitals are needed
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Klein, Organic Chemistry 2e
1-46Slide47
1.9 Hybridized Atomic Orbitals
The sp atomic orbitals overlap HEAD-ON to form sigma (σ) bonds while the unhybridized p orbitals overlap SIDE-BY-SIDE to form pi bonds
Practice with SkillBuilder 1.7Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e 1-
47Slide48
1.9 Hybridized Atomic Orbitals
Which should be stronger, a pi bond or a sigma bond? WHY?Which should be longer, an sp3 – sp3 sigma bond overlap or an sp
– sp sigma bond overlap?Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-48Slide49
1.9 Hybridized Atomic Orbitals
Explain the different strengths and lengths below.Practice with conceptual checkpoint 1.24
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1.10 Molecular Geometry
Valence shell electron pair repulsion (VSEPR theory)Valence electrons (bonded and lone pairs) repel each otherTo determine molecular geometry…Determine the Steric number
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-50Slide51
1.10 Molecular Geometry
Valence shell electron pair repulsion (VSEPR theory)Valence electrons (bonded and lone pairs) repel each otherTo determine molecular geometry…Predict the hybridization of the central atom If the Steric number is 4, then it is sp3
If the Steric number is 3, then it is sp2If the Steric number is 2, then it is sp
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Klein, Organic Chemistry 2e 1-
51Slide52
1.10 sp3 Geometry
For any sp3 hybridized atom, the 4 valence electron pairs will form a tetrahedral electron group geometry
Methane has 4 equal bonds, so the bond angles are equalHow does the lone pair of ammonia affect its geometry?
The bond angels in oxygen are even smaller, why?
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Klein, Organic Chemistry 2e
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1.10 sp3 Geometry
The molecular geometry is different from the electron group geometry. HOW?
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1.10 sp2 Geometry
Calculate the Steric number for BF3 Electron pairs that are located in sp2 hybridized orbitals will form a trigonal
planar electron group geometryWhat will be the molecular geometry?Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
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1.10 sp2 Geometry
How many electrons are in Boron’s unhybridized p orbital? Does this geometry follow VSEPR theory?Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e
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1.10 sp2 Geometry
Analyze the steric number, hybridization, electron group geometry and molecular geometry for this imine?
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-56Slide57
1.10 sp Geometry
Analyze the Steric number, the hybridization, the electron group geometry, and the molecular geometry for the following moleculesBeH2CO2
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e 1-57Slide58
1.10 Geometry Summary
Practice with SkillBuilder 1.8Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e 1-58Slide59
1.11 Molecular Polarity
Electronegativity Differences cause inductionInduction (shifting of electrons WITHIN their orbitals) results in a dipole moment.Dipole moment = (the amount of partial charge) x (the distance the δ+ and δ- are separated)Dipole moments are reported in units of
debye (D)1 debye = 10-18 esu ∙ cmAn esu is a unit of charge. 1 e- has a charge of 4.80 x 10-10
esucm are included in the unit, because the distance between the centers of + and – charges affects the dipoleCopyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e
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1.11 Molecular Polarity
Consider the dipole for CH3ClDipole moment (μ) = charge (e) x distance (d
)Plug in the charge and distanceμ = (1.056 x 10-10 esu) x (1.772 x 10-8 cm)Note that the amount of charge separation is less than what it would be if it were a full charge separation (4.80 x 10-10 esu)μ = 1.87 x 10
-18 esu ∙ cmConvert to debyeμ = 1.87 DCopyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e
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1.11 Molecular Polarity
What would the dipole moment be if CH3Cl were 100% ionic?μ = charge (e) x distance (d)
Plug in the charge and distanceμ = (4.80 x 10-10 esu) x (1.772 x 10-8 cm)The full charge of an electron is plugged inμ = 8.51 x 10-18 esu ∙ cm = 8.51 DWhat % of the C-Cl bond is ionic?
Is the C-Cl bond mostly ionic or mostly covalent?Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e
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1.11 Molecular Polarity
Check out the polarity of some other common bonds
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-62Slide63
1.11 Molecular Polarity
Why is the C=O double bond so much more polar than the C-O single bond?Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-63Slide64
1.11 Molecular Polarity
For molecules with multiple polar bonds, the dipole moment is the vector sum of all of the individual bond dipoles
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-64Slide65
1.11 Molecular Polarity
It is important to determine a molecule’s geometry FIRST before analyzing its polarityIf you have not drawn the molecule with the proper geometry, it may cause you to assess the polarity wrong as wellWould the dipole for water be different if it were linear rather than angular?Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e 1-65Slide66
1.11 Molecular Polarity
Electrostatic potential maps are often used to give a visual depiction of polarityCopyright © 2015 John Wiley & Sons, Inc. All rights reserved.
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1.11 Molecular Polarity
Practice with SkillBuilder 1.9Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e 1-67Slide68
1.11 Molecular Polarity
Explain why the dipole moment for pentane = 0 DCopyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-68Slide69
1.12 Intermolecular Forces
Many properties such as solubility, boiling point, density, state of matter, melting point, etc. are affected by the attractions BETWEEN moleculesNeutral molecules (polar and nonpolar) are attracted to one another through…Dipole-dipole interactionsHydrogen bondingDispersion forces (a.k.a. London forces or fleeting dipole-dipole forces)
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e 1-69Slide70
1.12 Dipole-Dipole
Dipole-dipole forces result when polar molecules line up their opposite charges.Note acetone’s permanent dipole results from the difference in electronegativity between C and OThe dipole-dipole attractions BETWEEN acetone molecules affects acetone’s boiling and melting points. HOW?
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-70Slide71
1.12 Dipole-Dipole
Why do isobutylene and acetone have such different MP and BPs?Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e 1-71Slide72
1.12 Hydrogen Bonding
Hydrogen bonds are an especially strong type of dipole-dipole attractionHydrogen bonds are strong because the partial + and – charges are relatively largeWhy are the partial charges in the H-bonding examples below relatively large?
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e 1-72Slide73
1.12 Hydrogen Bonding
Only when a hydrogen shares electrons with a highly electronegative atom (O, N, F) will it carry a large partial positive chargeThe large δ+ on the H atom can attract large δ– charges on other moleculesEven with the large partial charges, H-bonds are still about 20 times weaker than covalent bondsCompounds with H atoms that are capable of forming H-bonds are called protic
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-73Slide74
1.12 Hydrogen Bonding
Which of the following solvents are protic (capable of H-bonding), and which are not?Acetic acidDiethyl etherMethylene chloride (CH2
Cl2)Dimethyl sulfoxide
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Klein, Organic Chemistry 2e
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1.12 Hydrogen Bonding
Explain why the following isomers have different boiling pointsCopyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e 1-75Slide76
1.12 Hydrogen Bonding
H-bonds are among the forces that cause DNA to form a double helix and some proteins to fold into an alpha-helixCopyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e 1-76Slide77
1.12 London Dispersion Forces
If two molecules are nonpolar (dipole = 0 D), will they attract one another?YES! HOW?Nonpolar molecules normally have their electrons (–) spread out evenly around the nuclei (+) completely balancing the charge
However, the electrons are in constant random motion within their MOsCopyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e 1-
77Slide78
1.12 London Dispersion Forces
The constant random motion of the electrons in the molecule will sometimes produce an electron distribution that is NOT evenly balanced with the positive charge of the nucleiSuch uneven distribution produces a temporary dipole, which can induce a temporary dipole in a neighboring molecule
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e 1-78Slide79
1.12 London Dispersion Forces
The result is a fleeting attraction between the two moleculesSuch fleeting attractions are generally weak. But like any weak attraction, if there are enough of them, they can add up to a lot
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-79Slide80
1.12 London Dispersion Forces
The greater the surface area of a molecule, the more temporary dipole attractions are possibleConsider the feet of Gecko. They have many flexible hairs on their feet that maximize surface contact
The resulting London dispersion forces are strong enough to support the weight of the GeckoCopyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-80Slide81
1.12 London Dispersion Forces
Explain why molecules with more mass generally have higher boiling pointsPractice with SkillBuilder 1.10
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e 1-
81Slide82
1.12 London Dispersion Forces
Explain why more highly branched molecules generally have lower boiling pointsCopyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e 1-82Slide83
1.13 Solubility
We use the principle, like-dissolves-likePolar compounds generally mix well with other polar compoundsIf the compounds mixing are all capable of H-bonding and/or strong dipole-dipole, then there is no reason why they shouldn’t mixNonpolar compounds generally mix well with other nonpolar compoundsIf none of the compounds are capable of forming strong attractions, then no strong attractions would have to be broken to allow them to mix
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e 1-83Slide84
1.13 Solubility
We know it is difficult to get a polar compound (like water) to mix with a nonpolar compound (like oil)We can’t use just water to wash oil off our dirty clothsTo remove nonpolar oils, grease, and dirt, we need soap
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e 1-84Slide85
1.13 Solubility
Soap molecules organize into micelles in water, which form a nonpolar interior to carry away dirt.Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e 1-85Slide86
1.13 Solubility
Which attraction is generally stronger?The attraction between a permanent dipole and an induced dipole versusThe attraction between a temporary dipole and an induced dipoleWhich attraction is generally stronger?The attraction between a polar molecule and a nonpolar molecule
versusThe attraction between two nonpolar molecules?Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-86Slide87
1.13 Solubility
Why won’t a nonpolar compound readily dissolve in water?Is it because the water molecules repel the nonpolar molecules?Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-87Slide88
Additional Example Problems
Draw the structure for C2Cl3NDraw the structure for CH2O
Draw the structure for CH2O2Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.Klein, Organic Chemistry 2e
1-88Slide89
Additional Example Problems
Give all formal charges in the structures below.
Copyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e
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Additional Example Problems
How many nodes are in the 3s subshell? How many nodes are in a typical sigma antibonding MO? How many nodes are in a typical sp3 orbital?
How many nodes are in a typical pi bonding MO? How many nodes are in a typical pi antibonding MOCopyright © 2015 John Wiley & Sons, Inc. All rights reserved.
Klein, Organic Chemistry 2e 1-90Slide91
Additional Example Problems
Analyze the geometry, polarity and types of intermolecular attractions for the following molecules.
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Klein, Organic Chemistry 2e
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91