Organic chemistry lecture powerpoint BY WADE

Organic Chemistry, 8th Edition L. G. Wade, Jr.

Chapter 1 Lecture Introduction and Review

Rizalia Klausmeyer Baylor University Waco, TX © 2013 Pearson Education, Inc. © 2013 Pearson Education, Inc.

Organic Chemistry O H H CH3CH2CH CHCH2CONH O

2-Pentenylpenicillin

N

S

O

CH3

O

CH3

H5C6

COONa

O OH

C6H5 O N H

O OH

HO H5C6

CH3O

O O

O O

O

O

CH3

H NCH3

Taxol

HO

Codeine

• Organic chemistry is the chemistry of carbon compounds. © 2013 Pearson Education, Inc.

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Electronic Structure of the Atom • An atom has a dense, positively charged nucleus surrounded by a cloud of electrons. • The electron density is highest at the nucleus and drops off exponentially with increasing distance from the nucleus in any direction.

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The 2p Orbitals • There are three 2p orbitals, oriented at right angles to each other. • Each p orbital consists of two lobes. • Each is labeled according to its orientation along the x, y, or z axis.

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Isotopes 12

C 6

14

C 6

• Isotopes are atoms with the same number of protons but different number of neutrons. • Mass number is the sum of the protons and neutrons in an atom. © 2013 Pearson Education, Inc.

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Electronic Configurations of Atoms • Valence electrons are electrons on the outermost shell of the atom.

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Electronic Configurations • The aufbau principle states to fill the lowest energy orbitals first. • Hund’s rule states that when there are two or more orbitals of the same energy (degenerate), electrons will go into different orbitals rather than pairing up in the same orbital.

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Ionic Bonding • To obtain a noble gas configuration (a full valence shell), atoms may transfer electrons from one atom to another. • The atoms, now bearing opposite charges, attract each other, forming an ionic bond.

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Covalent Bonding • Electrons are shared between the atoms to complete the octet. • When the electrons are shared evenly, the bond is said to be nonpolar covalent, or pure covalent. • When electrons are not shared evenly between the atoms, the resulting bond will be polar covalent.

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Lewis Structures CH4

NH3 H

Carbon: 4 e 4 H@1 e ea: 4 e 8e

H C H

Nitrogen: 5 e 3 H@1 e ea: 3 e 8e

H N H H

H

H2O Oxygen: 6 e 2 H@1 e ea: 2 e 8e © 2013 Pearson Education, Inc.

Cl2 H O H

2 Cl @7 e ea: 14 e

Chapter 1

Cl Cl 10

Bonding Patterns Valence electrons

C N

O Halides (F, Cl, Br, I) © 2013 Pearson Education, Inc.

# Bonds

# Lone Pair Electrons

4

4

0

5

3

1

6

2

2

7

1

3

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Nonbonding Electrons

• Also called lone pairs. • Nonbonding electrons are valence shell electrons that are not shared between atoms.

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Multiple Bonding

• Sharing two pairs of electrons is called a double bond. • Sharing three pairs of electrons is called a triple bond. © 2013 Pearson Education, Inc.

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Lewis structures are the way we write organic chemistry. Learning now to draw them quickly and correctly will help you throughout this course.

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Dipole Moment

• Dipole moment is defined to be the amount of charge separation (d) multiplied by the bond length (m). • Charge separation is shown by an electrostatic potential map (EPM), where red indicates a partially negative region and blue indicates a partially positive region. © 2013 Pearson Education, Inc.

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Pauling Electronegativities • Electronegativities can be used to predict whether a bond will be polar and the direction of its dipole moment. • Since the electronegativities of carbon and hydrogen are similar, C—H bonds are considered to be nonpolar.

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Formal Charges Formal charge = [group number ] – [nonbonding electrons ] – ½ [shared electrons]

H3O+

NO+ 6 – 2 – ½ (6) = +1

6 – 2 – ½ (6) = +1 +

N O

H O H H

5 – 2 – ½ (6) = 0

+

• Formal charges are a way of keeping track of electrons. • They may or may not correspond to actual charges in the molecule.

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Solved Problem 1 Compute the formal charge (FC) on each atom in H3N – BH3.

Solution

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Common Bonding Patterns

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Work enough problems to become familiar with these bonding patterns so you can recognize other patterns as being either unusual or wrong.

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Resonance Forms • The structures of some compounds are not adequately represented by a single Lewis structure. • Resonance forms are Lewis structures that can be interconverted by moving electrons only. • The true structure will be a hybrid between the contributing resonance forms.

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Resonance Forms Resonance forms can be compared using the following criteria, beginning with the most important: 1. Has as many octets as possible. 2. Has as many bonds as possible. 3. Has the negative charge on the most electronegative atom. 4. Has as little charge separation as possible.

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Major and Minor Contributors • The major contributor is the one in which all the atoms have a complete octet of electrons.

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Major and Minor Contributors (Continued) • When both resonance forms obey the octet rule, the major contributor is the one with the negative charge on the most electronegative atom.

N C O

N C O

MAJOR

MINOR

The oxygen is more electronegative, so it should have the negative charge. © 2013 Pearson Education, Inc.

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Non-Equivalent Resonance • Opposite charges should be on adjacent atoms. CH3 O N O

CH3 O N O

MAJOR

MINOR

The most stable one is the one with the smallest separation of oppositely charged atoms.

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Solved Problem 2 Draw the important resonance forms for [CH3OCH2]+. Indicate which structure is the major and minor contributor or whether they would have the same energy.

Solution

The first (minor) structure has a carbon atom with only six electrons around it. The second (major) structure has octets on all atoms and an additional bond.

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Solved Problem 3 Draw the resonance structures of the compound below. Indicate which structure is the major and minor contributor or whether they would have the same energy.

Solution

Both of these structures have octets on oxygen and both carbon atoms, and they have the same number of bonds. The first structure has the negative charge on carbon, the second on oxygen. Oxygen is the more electronegative element, so the second structure is the major contributor.

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Resonance Forms for the Acetate Ion

• When acetic acid loses a proton, the resulting acetate ion has a negative charge delocalized over both oxygen atoms. • Each oxygen atom bears half of the negative charge, and this delocalization stabilizes the ion. • Each of the carbon–oxygen bonds is halfway between a single bond and a double bond and is said to have a bond order of 1½. © 2013 Pearson Education, Inc.

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Condensed Structural Formulas Lewis

Condensed

H H 1

2

CH3CH3

H C C H H H

• Condensed forms are written without showing all the individual bonds. • Atoms bonded to the central atom are listed after the central atom (CH3CH3, not H3CCH3). • If there are two or more identical groups, parentheses and a subscript may be used to represent them. © 2013 Pearson Education, Inc.

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Condensed Structural Formulas Lewis

Condensed Br

H H Br H 1

2

3

H C C C

4

C H

CH3CH2CHCH3

H H H H

O

H H O H 1

2

3

H C C C OH H © 2013 Pearson Education, Inc.

4

HOCH2CH2CCH3

C H

H Chapter 1

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Condensed Structural Formulas (Continued)

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Condensed Structural Formulas (Continued)

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Condensed Structural Formulas (Continued)

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Line-Angle Drawings H H H H 1

2

H C C C

3

4

C H

C

C

C

C

H H H H • Sometimes called skeletal structure or stick figure. • Bonds are represented by lines, and carbons are present where a line begins or ends and where two lines meet. • Hydrogens attached to carbon are not shown. • Nitrogen, oxygen, and halides must be shown.

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Line-Angle Drawings H H H H H 1

2

3

H C C C

4

O

1

2

3

4

5

6

5 6

O

C C C

H H H H H

H

H

• Atoms other than carbon must be shown. • Double and triple bonds must also be shown.

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Line-Angle Formulas (Continued)

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Calculating Empirical Formulas The following are items that need to be considered when calculating empirical formulas: – Given % composition for each element, assume 100 g. – Convert the grams of each element to moles. – Divide by the smallest number of moles to get ratio. – Molecular formula may be a multiple of the empirical formula. © 2013 Pearson Education, Inc.

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If an elemental analysis does not add up to 100%, the missing percentage is assumed to be oxygen.

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Arrhenius Acids • Arrhenius acids are substances that dissociate in water to give H3O+ ions.

• Stronger acids dissociate to a greater degree than weaker acids. © 2013 Pearson Education, Inc.

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Arrhenius Bases • Arrhenius bases are substances that dissociate in water to give hydroxide ions.

• Stronger bases (NaOH) dissociate more than weaker bases (Mg(OH)2). © 2013 Pearson Education, Inc.

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Brønsted-Lowry Acids and Bases Brønsted-Lowry acids are any species that donate a proton. Brønsted-Lowry bases are any species that can accept a proton.

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Conjugate Acids and Bases

• Conjugate acid: when a base accepts a proton, it becomes an acid capable of returning that proton. • Conjugate base: when an acid donates its proton, it becomes capable of accepting that proton back. © 2013 Pearson Education, Inc.

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Effect of Electronegativity on pKa • As the bond to H becomes more polarized, H becomes more positive and the bond is easier to break.

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Effect of Size on pKa

• As size increases, the H is more loosely held and the bond is easier to break. • A larger size also stabilizes the anion. © 2013 Pearson Education, Inc.

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Effect of Resonance on pKa

• If the negative charge on an atom can be delocalized over two or more atoms, the acidity of that compound will be greater than when the negative charge cannot be delocalized. • The ethoxide anion is less acidic than the acetate ion simply because the acetate ion can delocalize the negative charge. • Methanesulfonic acid can delocalize the charge in three different resonance forms, making it more acidic than the acetate ion. © 2013 Pearson Education, Inc.

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Lewis Acids and Lewis Bases • Lewis bases are species with available electrons than can be donated to form a new bond. • Lewis acids are species that can accept these electrons to form new bonds. • Since a Lewis acid accepts a pair of electrons, it is called an electrophile.

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Nucleophiles and Electrophiles • Nucleophile: Donates electrons to a nucleus with an empty orbital. • Electrophile: Accepts a pair of electrons. • When forming a bond, the nucleophile attacks the electrophile, so the arrow goes from negative to positive. • When breaking a bond, the more electronegative atom receives the electrons. © 2013 Pearson Education, Inc.

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Nucleophiles and Electrophiles

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