Klein Chapter 1

1.1 Organic Chemistry• The study of carbon-containing molecules and their

reactions• What happens to a molecule during a reaction?

– A collision– Bonds break/form

• WHAT IS A BOND?

• The BIG question: WHY do reactions occur?– We will need at least 2 semesters of your time to answer

this question– FOCUS ON THE ELECTRONS

Klein, Organic Chemistry 1e Copyright 2012 John Wiley & Sons, Inc. 1-1

1.1 Organic Chemistry• Why do we distinguish between organic and

inorganic compounds?

• Why are organic compounds important?


1.2 Structural Theory• In the mid 1800s, it was first suggested that

substances are defined by a specific arrangement of atoms.– Why is a compound’s formula NOT adequate to define it?

• What term do we use to describe different substances with the same formula?


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 atoms

• Practice with skill builder 1.1


1.3 Covalent Bonding• A covalent bond is a PAIR of electrons shared between

two atoms. For example…


1.3 Covalent Bonding• How do potential energy and stability relate?

• What forces keep the bond at the optimal length?


1.3 Atomic Structure• A review from General Chemistry

– Protons (+1) and neutrons (neutral) reside in the nucleus– Electrons (-1) reside outside the nucleus. WHERE?

– Some electrons are close to the nucleus and others are far away, WHY?

– Look at carbon for example. Which electrons are the valence electrons?

– Why are valence electrons important?


• You can always calculate the number of valence electron by analyzing the e- configuration. Look at phosphorus.

• Or, for Group A elements only, just look at the Group number (Roman Numeral) on the periodic table


1.3 Counting Valence Electrons


1.3 Simple Lewis Structures• For simple Lewis Structures…

1. Draw the individual atoms using dots to represent the valence electrons.

2. Put the atoms together so they share PAIRS of electrons to make complete octets. WHAT is an octet?

• Take NH3, for example…



1.3 Simple Lewis Structures• For simple Lewis Structures…

1. Draw the individual atoms using dots to represent the valence electrons.

2. Put the atoms together so they share PAIRS of electrons to make complete octets. WHAT is an octet?

• Try drawing the structure for C2H2


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 stability

• To 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


1.4 Formal Charge• Lets look at a specific formal charge example. Given the

Lewis structure, calculate the formal charge on each atom.

or

• Carbon should have 4 valence e-s, because it is in group IVA on the periodic table.

• Carbon actually has 8 valence e-s. 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.


1.4 Formal Charge• Lets look at the formal charge of the oxygen atom.

or

• Oxygen should have 6 valence e-s, because it is in group VIA on the periodic table.

• It actually has 8 valence e-s. 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 skill builder 1.4Klein, Organic Chemistry 1e Copyright 2012 John Wiley & Sons, Inc. 1-13

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?


1.5 Polar Covalent Bonds• Covalent bonds are either polar or nonpolar

– Nonpolar Covalent –bonded atoms share electrons evenly– Polar Covalent – One of the atoms attracts electrons more than

the other

• Electronegativty - how strongly an atom attracts shared electrons


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.


1.5 Polar Covalent Bonds• Can a bond have both covalent and ionic character?



1.6 Atomic OrbitalsGeneral Chemistry review• In the 1920s, Quantum Mechanics was established as a

theory to explain the wave properties of electrons• The solution to wave equations for electrons provides us

with visual pictures called orbitals


1.6 Atomic OrbitalsGeneral Chemistry review• The type or orbital be identified by its shape• An 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


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.


1.6 Atomic Orbitals• Because they are generated

mathematically from wavefunctions, orbital regions can also be (-), (+), or ZERO– The 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.


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?


• Once the 2s is full, electrons fill into the three degenerate 2p orbitals

• Where are the nodes in each of the 2p orbitals?

1.6 Atomic Orbitals


1.6 Atomic Orbitals• Common elements and their electron configurations



1.6 Atomic Orbitals• What are the rules that govern our placement of

electrons ?


1.7 Valence Bond Theory• A bond occurs when atomic orbitals overlap. Overlapping

orbitals is like overlapping waves

• Only constructive interference results in a bond


1.7 Valence Bond Theory• The bond for a H2 molecule results from constructive interference

• Where do the bonded electrons spend most of their time?


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 MOs

• 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


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 it MOs occupied, will be stable or unstable?


• How would the energy of the He2 compare to 2 He?

• Why does Helium exist in its atomic form rather than in molecular form?

1.8 Molecular Orbital Theory• Consider the MOs for CHBr3

– There are many areas of atomic orbital overlap– Notice how the MOs extend over the entire molecule

• Each picture below represents ONE orbital.


1.8 Molecular Orbital Theory• How many electrons can fit into the areas represented in

(b)?

• Depending on the circumstances, we will use both MO and valence bond theory to explain phenomena.


1.9 Hybridized Atomic Orbitals• Given the electron configuration for C and H, imagine how their atomic orbitals

might overlap

• Would such orbital overlap yield methane?


1.9 Hybridized Atomic Orbitals• To make methane, the C atom must have 4 atomic orbitals available for overlapping• If 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?


1.9 Hybridized Atomic Orbitals• The carbon must undergo hybridization to form 4 equal

atomic orbitals• The atomic orbitals must be equal in energy to form four

equal-energy symmetrical C-H bonds


1.9 Hybridized Atomic Orbitals

• Should the shape of an sp3 orbital look more like an s or more like p orbital?


1.9 Hybridized Atomic Orbitals• To make CH4, the 1s atomic orbitals of the H atoms will

overlap with the four sp3 hybrid atomic orbitals of C


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


1.9 Hybridized Atomic Orbitals• Consider ethene (ethylene).

• Each carbon in ethene must bond to three other atoms, so only three hybridized atomic orbitals are needed



• An sp2 hybridized carbon will have three equal-energy sp2 orbitals and one unhybridized p orbital


1.9 Hybridized Atomic Orbitals• The sp2 atomic orbitals overlap to form sigma (σ) bonds

• Sigma bonds provide maximum HEAD-ON overlap


1.9 Hybridized Atomic Orbitals• The unhybridized p orbitals in ethene form pi (π) bonds, SIDE-BY-SIDE overlap



1.9 Hybridized Atomic Orbitals• The unhybridized p orbitals in ethene form pi (π) bonds,

SIDE-BY-SIDE overlap


• MO theory provides a similar picture. Remember, red and blue regions are all part of the same orbital



• Why is sp2 hybridization not appropriate for methane (CH4)?

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


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



1.9 Hybridized Atomic Orbitals• Which should be stronger, a pi bond or a sigma bond? WHY?

• Which should be stronger, an sp3 – sp3 sigma bond overlap or an sp – sp sigma bond overlap?


1.9 Hybridized Atomic Orbitals• Explain the different strengths and lengths below.



1.10 Molecular Geometry• Valence shell electron pair repulsion (VSEPR theory)

– Valence electrons (bonded and lone pairs) repel each other

• To determine molecular geometry…1. Determine the Steric number


1.10 Molecular Geometry• Valence shell electron pair repulsion (VSEPR theory)

– Valence electrons (bonded and lone pairs) repel each other

• To determine molecular geometry…2. 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 sp2

• If the Steric number is 2, then it is sp


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 angels are equal

• How does the lone pair of ammonia affect its geometry?

• The bond angels in oxygen are even smaller, why?

1.10 sp3 Geometry• The molecular geometry is different from the electron

group geometry. HOW?


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 geometry

• What will be the molecular geometry?


1.10 sp2 Geometry


• How many electrons are in Boron’s unhybridized p orbital? • Does this geometry follow VSEPR theory?

1.10 sp2 Geometry


Analyze the steric number, hybridization, electron group geometry and molecular geometry for this imine?

1.10 sp Geometry• Analyze the Steric number, the hybridization, the electron group geometry, and the molecular geometry for the following molecules• BeH2

• CO2


1.10 Geometry Summary



1.11 Molecular Polarity• Electronegativity Differences cause induction• Induction (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 cm∙

– An esu is a unit of charge. 1 e- has a charge of 4.80 x 10-10 esu– cm are included in the unit, because the distance between the

centers of + and – charges affects the dipole


1.11 Molecular Polarity• Lets consider the dipole of CH3Cl

• Dipole moment (μ) = charge (e) x distance (d)


– Plug in the charge and distance

• μ = (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 cm∙– Convert to debye

• μ = 1.87 D

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 distance

• μ = (4.80 x 10-10 esu) x (1.772 x 10-8 cm)– The full charge of an electron is plugged in

• μ = 8.51 x 10-18 esu cm = ∙ 8.51 D• What % of the C-Cl bond is ionic?• Is the C-Cl bond mostly ionic or mostly

covalent?

1.11 Molecular Polarity• Check out the polarity of come other common bonds


1.11 Molecular Polarity• Why is the C=O double bond so much more polar than

the C-O single bond?


1.11 Molecular Polarity• For molecules with multiple polar bonds, the dipole

moment is the vector sum of all of the individual bond dipoles


1.11 Molecular Polarity• It is important to determine a molecule’s geometry FIRST before analyzing its polarity• If you have not drawn the molecule with the proper geometry, it may cause you to

aasses the polarity wrong as well• Would the dipole for water be different if it were linear rather than angular?


1.11 Molecular Polarity



1.11 Molecular Polarity

• Explain why the dipole moment for pentane = 0 D


1.12 Intermolecular Forces• Many properties such as solubility, boiling point, density, state of matter, melting point, etc. are affected by

the attractions BETWEEN molecules• Neutral molecules (polar and nonpolar) are attracted to one another through…

– Dipole-dipole interactions– Hydrogen bonding– Dispersion forces (a.k.a. London forces or fleeting dipole-dipole forces)


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 O• The dipole-dipole attractions BETWEEN acetone molecules affects acetone’s boiling and

melting points. HOW?


1.12 Dipole-Dipole


• Why do isobutylene and acetone have such different MP and BPs?

1.12 Hydrogen Bonding• Hydrogen bonds are an especially strong type of dipole-

dipole attraction• Hydrogen bonds are strong because the partial + and –

charges are relatively large• Why are the partial charges in the H-bonding examples

below relatively large?


1.12 Hydrogen Bonding• Only when a hydrogen shares electrons with a highly electronegative

atom (O, N, F, or Cl) will it carry a large partial positive charge• The large δ+ on the H atom can attract large δ– charges on other

molecules• Even with the large partial charges, H-bonds are still about 20 times

weaker than covalent bonds• Compounds with H atoms that are capable of forming H-bonds are

called protic


1.12 Hydrogen Bonding• Which of the following solvents are protic (capable of H-bonding), and which are not?• Acetic acid

• Diethyl ether

• Methylene chloride (CH2Cl2)

• Dimethyl sulfoxide


H3C

C

OH

O

H3C

H2C

O

H2C

CH3

H3C

S

O

CH3

1.12 Hydrogen Bonding• Explain why the following isomers have different boiling

points


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 MOs


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 nuclei

• Such uneven distribution produces a temporary dipole, which can induce a temporary dipole in a neighboring molecule


1.12 London Dispersion Forces• The result is a fleeting attraction between the two molecules

• Such fleeting attractions are generally weak. • But like any weak attraction, if there are enough of them, they can add up to a

lot


1.12 London Dispersion Forces• The greater the surface area of a molecule, the more

temporary dipole attractions are possible• Consider 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 Gecko

1.12 London Dispersion Forces• Explain why molecules with more mass generally have higher boiling points



1.12 London Dispersion Forces• Explain why more highly branched molecules generally

have lower boiling points


1.13 Solubility• We use the principle, like-dissolves-like• Polar compounds generally mix well with other polar compounds

– If the compounds mixing are all capable of H-bonding and/or strong dipole-dipole, then there is no reason why they shouldn’t mix

• Nonpolar compounds generally mix well with other nonpolar compounds– If none of the compounds are capable of forming strong attractions,

then no strong attractions would have to be broken to allow them to mix


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 cloths

• To remove nonpolar oils, grease, and dirt, we need soap


1.13 Solubility• Soap molecules organize into micelles in water, which

form a nonpolar interior to carry away dirt.


1.13 Solubility• Which attraction is generally stronger?

– The attraction between a permanent dipole and an induced dipoleor

– The attraction between a temporary dipole and an induced dipole

• Which attraction is generally stronger?– The attraction between a polar molecule and a nonpolar molecule

or– The attraction between two nonpolar molecules?


1.13 Solubility• Why won’t a nonpolar compound readily dissolve in

water?• Is it because the water molecules repel the nonpolar

molecules?


Klein Chapter 1

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