Organic Chemistry: Fundamental Concepts and Reactions, Summaries of Law

Download Organic Chemistry: Fundamental Concepts and Reactions and more Summaries Law in PDF only on Docsity!

ORGANIC CHEMISTRY I

Xin Liu

Kwantlen Polytechnic University

Kwantlen Polytechnic University

Organic Chemistry I

Xin Liu

1 https://chem.libretexts.org/@go/page/

TABLE OF CONTENTS

Licensing

About the Author

Acknowledgements

Introduction

1: Basic Concepts in Chemical Bonding and Organic Molecules

1.1: Chemical Bonding 1.2: Lewis Structure 1.3: Resonance Structures 1.4: Resonance structures in Organic Chemistry 1.5: Valence-Shell Electron-Pair Repulsion Theory (VSEPR) 1.6: Valence Bond Theory and Hybridization 1.7: Answers to Practice Questions Chapter 1

2: Fundamental of Organic Structures

2.1: Structures of Alkenes 2.2: Nomenclature of Alkanes 2.3: Functional Groups 2.4: IUPAC Naming of Organic Compounds with Functional Groups 2.5: Degree of Unsaturation/Index of Hydrogen Deficiency 2.6: Intermolecular Force and Physical Properties of Organic Compounds 2.7: Answers to Practice Questions Chapter 2

3: Acids and Bases- Organic Reaction Mechanism Introduction

3.1: Review of Acids and Bases and Ka 3.2: Organic Acids and Bases and Organic Reaction Mechanism 3.3: pKa of Organic Acids and Application of pKa to Predict Acid-Base Reaction Outcome 3.4: Structural Effects on Acidity and Basicity 3.5: Lewis Acids and Lewis Bases 3.6: Answers to Practice Questions Chapter 3

4: Conformations of Alkanes and Cycloalkanes

4.1: Conformation Analysis of Alkanes 4.2: Cycloalkanes and Their Relative Stabilities 4.3: Conformation Analysis of Cyclohexane 4.4: Substituted Cyclohexanes 4.5: Answers to Practice Questions Chapter 4

5: Stereochemistry

5.1: Summary of Isomers 5.2: Geometric Isomers and E/Z Naming System 5.3: Chirality and R/S Naming System 5.4: Optical Activity

2 https://chem.libretexts.org/@go/page/

5.5: Fisher Projection 5.6: Compounds with More Than One Chirality Centers 5.7: Answers to Practice Questions Chapter 5

6: Structural Identification of Organic Compounds- IR and NMR

Spectroscopy

6.1: Electromagnetic Radiation and Molecular Spectroscopy 6.2: Infrared (IR) Spectroscopy Theory 6.3: IR Spectrum and Characteristic Absorption Bands 6.4: IR Spectrum Interpretation Practice 6.5: NMR Theory and Experiment 6.6: ¹H NMR Spectra and Interpretation (Part I) 6.7: ¹H NMR Spectra and Interpretation (Part II) 6.8: ¹³C NMR Spectroscopy 6.9: Structure Determination Practice 6.10: Answers to Practice Questions Chapter 6

7: Nucleophilic Substitution Reactions

7.1: Nucleophilic Substitution Reaction Overview 7.2: SN2 Reaction Mechanism, Energy Diagram and Stereochemistry 7.3: Other Factors that Affect SN2 Reactions 7.4: SN1 Reaction Mechanism, Energy Diagram and Stereochemistry 7.5: SN1 vs SN 7.6: Extra Topics on Nucleophilic Substitution Reaction 7.7: Answers to Practice Questions Chapter 7

8: Elimination Reactions

8.1: E2 Reaction 8.2: E1 Reaction 8.3: E1/E2 Summary 8.4: Comparison and Competition Between SN1, SN2, E1 and E 8.5: Answers to Practice Questions Chapter 8

9: Free Radical Substitution Reaction of Alkanes

9.1: Homolytic and Heterolytic Cleavage 9.2: Halogenation Reaction of Alkanes 9.3: Stability of Alkyl Radicals 9.4: Chlorination vs Bromination 9.5: Stereochemistry for Halogenation of Alkanes 9.6: Synthesis of Target Molecules- Introduction of Retrosynthetic Analysis 9.7: Answers to Practice Questions Chapter 9

10: Alkenes and Alkynes

10.1: Synthesis of Alkenes 10.2: Reactions of Alkenes- Addition of Hydrogen Halide to Alkenes 10.3: Reactions of Alkenes- Addition of Water (or Alcohol) to Alkenes 10.4: Reactions of Alkenes- Addition of Bromine and Chlorine to Alkenes 10.5: Reaction of Alkenes- Hydrogenation 10.6: Two Other Hydration Reactions of Alkenes

1 https://chem.libretexts.org/@go/page/

Licensing

A detailed breakdown of this resource's licensing can be found in Back Matter/Detailed Licensing.

1 https://chem.libretexts.org/@go/page/

About the Author

Xin Liu, Kwantlen Polytechnic University

xin.liu@kpu.ca

Dr. Xin Liu, Author

Dr. Xin Liu has been a faculty member at the Department of Chemistry, KPU since 2008. Other than teaching Organic Chemistry and first-year General Chemistry courses, she has also been actively involved in curriculum review and new course development. Having a keen interest in Open Education Resource, Dr. Liu hope to make more contributions in this fast growing area, to make learning accessible to everyone.

1 https://chem.libretexts.org/@go/page/

Introduction

What Is Organic Chemistry and Why Is It Important?

On a lovely Saturday afternoon in April, you are relaxing in a garden whilst enjoying a hot cup of coffee. Colourful spring blossoms lace the air with a pleasant aroma, and the green grass, warm sunshine and rich espresso make the afternoon a charming occasion.

Your mind begins to drift as you contemplate the combination of scents, colours and tastes that surround you in this moment, and how they make up the human experience’s unique and fascinating complexities.

If you have ever wondered about the origins of nature’s vibrant hues or the reasonings behind the alluring flavour of coffee, you would be able to find every answer within the elaborate spectrum of knowledge in the study of Organic Chemistry.

Organic chemistry is the chemistry of compounds containing the carbon element: the common element of all living organisms. Anthocyanins are the pigments that give flowers their various colours, chlorophyll is responsible for the green shades of grass is involved in the photosynthesis process of plants, and caffeine is what makes coffee function the way that it does. All these substances contain carbon , and they are all organic compounds.

Image of a Cyanidin, Pelargondin, Chlorophyll a, and Caffeine

The root of the term organic dates back from over two hundred years ago, when its original meaning did not even involve the element of carbon. The word organic was first introduced in 1807 by Jöns Jakob Berzelius, a Swedish chemist, and was used to refer to compounds derived from living organisms. It was once believed that organic compounds could only be obtained directly from nature as they contained a mystical essence of life known as “vita force”, therefore making it impossible to create organic compounds artificially. This theory was shattered by a famous experiment conducted by German chemist Friedrich Wohler in 1828. In his experiment, Wohler successfully synthesized the crystal urea by heating ammonia and cyanic acid together. The synthesis of urea marked a new era in the history of organic chemistry, not only redefining the term organic , but also rerouting organic chemistry into a completely new scientific discipline. The contemporary definition of organic, being carbon-containing compounds, is now the scientific way of describing the term. However, it has remained true over the years that organic compounds are essential to every known lifeform, as an abundance of organic molecules constitute all living organisms.

There are two additional notes regarding the modern definition of organic. Firstly, while it is true that organic compounds are those containing the element carbon, it is important to know that not all compounds that contain carbon are organic compounds. For example, calcium carbonate (CaCO3), the primary component in certain rocks and chalk, can never be labelled as organic. Secondly, the “organic” food that is often found in supermarkets refers to the fact that the agricultural products were grown without the use of artificial pesticides, herbicides, or synthetic fertilizers, and has nothing to do with the presence of carbon in their chemical structures. This use of the word organic is possibly derived from the old definition, implying that the products came from nature, without human intervention.

2 https://chem.libretexts.org/@go/page/

As you may have been able to deduce, organic chemistry can be found in every corner of the world around us. From the food we eat, (the carbohydrates in bread, the protein in meat, the fructose in fruit, and more) to the fabric we wear, (cotton, nylon, polyester) and the fuels that power the technology around us (gasoline, natural gas, coal), the list of organic compounds involved in our lives is endless. An important significance in the application of organic chemistry is its critical role in the development of medicine and pharmaceuticals. The active ingredients found in medicine are most often organic compounds, either isolated from naturally occurring materials or synthesized in a lab. Just a few well-known examples include Aspirin, Tylenol, penicillin, insulin, Warfarin, and Tamiflu. The rapid developments of the pharmaceutical industry, in which organic chemistry has acted as a major driver, have saved millions of lives and has dramatically improved today’s quality of life.

The magic element that is the key to organic chemistry and all living organisms is carbon. What is it about the carbon element that makes it so special? This can mainly be attributed to the special bonding ability that carbon possesses. Carbon atoms can form strong covalent bonds with other carbon atoms in the form of chains and rings, and it also forms strong bonds with other elements such as hydrogen, oxygen, nitrogen, sulfur and more. As a result, the structures of organic compounds are hugely diverse and can be rather complex.

Tips for Studying Organic Chemistry

Learning organic chemistry can be both exciting and challenging. The most commonly misleading learning strategy is the notion that “I can be successful by simply memorizing everything”. While memorization may be necessary at times, it is but a small fraction of what is needed to learning organic chemistry; the more important factor is your understanding. There are many structures, reactions and mechanisms involved in the course, and surface-level memorization will not carry you all the way through. However, if you know the connections between the structures, understand the underlying principles of the reactivity of certain compounds, and can tell the similarities and differences between different mechanisms, you will find that it becomes much easier. A few suggestions for learning include:

Rewrite your own notes when studying. For example, restate the concepts in your own words, or write a map of the concepts that are related. Practice makes perfect. Do as many practice questions as you can, and try to make your own questions to double check your understanding. Use molecule model sets for certain topics.

About the Book

Due to the high price barrier, about half of Organic Chemistry students at KPU do not have access to the textbook. This has become a serious issue that significantly affects the learning outcomes for the course. The creation of this open textbook is intended to provide a solution to this problem and help students get success in this course.

The book contains ten chapters, with the contents cover from the basic concepts on chemical bonding, functional group, to stereochemistry, spectroscopy for structure determination (IR and NMR) and organic reactions (nucleophilic substitution, elimination, radical substitution of alkanes, addition and oxidation reactions of alkenes, preparation and reactions of alkynes).

Organic Chemistry is a challenging subject for many students. To help readers understand the concepts more easily, simple and concise languages are intentionally applied in the book. The featured shaded textbox areas are included frequently in the book, where readers can find useful learning tips, reminder of common errors, comparison between similar concepts. To help readers develop problem-solving skills, a small section labelled as “strategy” is usually given for the examples in the book. Readers are encouraged to try solving the problems by themselves with helpful hints provided in the “strategy”, and then compare their work with the detailed solutions provided afterwards. Because of limited time availability, not many practice questions are included in this book. We hope that more questions, particularly questions with different level of difficulties, will be added in the future edition of the book.

1.1.1 https://chem.libretexts.org/@go/page/

1.1: Chemical Bonding To summarize simply, a chemical bond is the attractive force holding atoms or ions together. Such attractive interaction leads to a more stable state for the whole system comparing to individual atoms. Valence electrons play a fundamental role in chemical bonding. In the electron configuration of an atom, the outermost shell is called the valence shell, and the electrons in the valence shell (outermost shell) are known as valence electrons. Take the carbon atom for example: the electron configuration of carbon is 1s 2s 2p. The outermost shell is the 2 principal shell, so there are 4 valence electrons in carbon. Valence electrons are the electrons that are the furthest away from the nucleus, and thus experience the least attraction from the nucleus and therefore are most reactive. They play the most important role in chemical bonding. Exercises 1. Determine the number of valence electrons for following elements: B, N, O, Cl, Mg. Answers to Practice Questions Chapter 1 Ionic Bond and Covalent Bond There are two major types of chemical bonding: ionic bonds and covalent bonds. An ionic bond is a bond that results from the electrostatic attraction (force) between ions of opposite charges. Ionic bonds apply to ionic compound, such as sodium chloride (NaCl). In simple ionic compounds, the metal element loses valence electron(s) to form the cation and the non-metal element gains electron(s) to form the anion. With the proper number of electron(s) lost or gained, both the cation and the anion achieve a full outer shell that contains eight electrons, as in the following examples of Na , Ca , Cl and O. According to Lewis’s Theory, an atom is most stable if its outer shell is filled or contains eight electrons. This is also called the octet rule. Na (atom) → Na + e Ca (atom) → Ca + 2e Cl (atom) + e → Cl O (atom) + 2e → O A covalent bond is a bond formed through the sharing of electron pairs between the two bonding atoms. The shared electron pairs are mutually attracted by the nuclei of both atoms. By sharing the electron pairs, both atoms also gain a filled outer shell, or an octet. Almost all of the bonds involved in organic compounds are covalent bonds. Covalent bond can be non-polar or polar. For covalent bonds formed between two identical atoms, the electron pairs are shared equally between the two nuclei. Electron density is distributed evenly through the bond, making the bond a non-polar bond. Examples include all homonuclear molecules, such as H-H, Cl-Cl, O=O, N≡N.

Figure 1.1a nonpolar covalent bond and polar covalent bond For heteronuclear bonds (the bond formed between two different atoms), the electron pairs are not shared evenly, and the bond is polar. The electron pairs are more attracted to the atom that has the stronger ability to pull the electron pairs towards itself. This ability is measured with electronegativity. The relative values of electronegativity (EN) are listed using the scale devised by Linus Pauling, as summarized in the following table:

Figure 1.1b Electronegativity Values in Pauling Scale

Notes about electronegativity values for Organic Chemistry purposes: It is much more important to know the trend of electronegativity than to memorize the values. The trend is that EN values decrease along the group from top to bottom and increase along the period from left to right (the trend mainly works for Main Group elements, not transition metal elements). It is very useful (although not mandatory) to know the EN values of a few select elements: F (4.0, highest), O (3.5), N (3.0), C (2.5) and H (2.1). The EN of C (2.5) and H (2.1) is rather close, which makes the C-H bond (the bond involved in all organic compounds) technically non-polar. With the introduction to the concept of electronegativity, bond polarity can be represented with the electronegativity difference between the two bonding atoms, which is known as ΔEN. For non- polar bonds, ΔEN equals to zero, and for polar bonds, ΔEN is not zero. The greater the ΔEN, the more polar the bond is. Exercises 1.

1. Identify the following bonds as “polar” or “non-polar”: C-C, C-H, B-F, O-O, C=N
2. Rank the following bonds in order of increasing bonding polarity: C—S, C—O, C—F (referring to the trend of EN, you do not need to use the exact EN values). Answers to Practice Questions Chapter 1 Because of the electronegativity difference, the atom with the higher EN attracts the shared electron pairs more strongly, therefore bearing a slightly negative charge (δ-). The other atom with a lower EN bears a slightly positive charge (δ+). The direction of the bond polarity can be indicated with an arrow, with the head of the arrow pointing to the negative end and a short perpendicular line near the tail of the arrow marking the positive end. The following example of an H-Cl molecule indicates how to show the bond polarity and partial charges of the polar bond.

2 2 2 nd

• 2 – 2–

• – 2+ –

• – – 2-

1.1.2 https://chem.libretexts.org/@go/page/

This page titled 1.1: Chemical Bonding is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by Xin Liu (Kwantlen Polytechnic University).

1.2.2 https://chem.libretexts.org/@go/page/

1. Calculate the total number of valence electrons. For ions, make sure charges are properly included in the calculation. For example of NH cation:

the total number of electrons = 5 (N atom) + 4×1 (four H atoms) -1 (minus the charge for cation) = 8 valence electrons

2. Write a plausible skeletal structure using the following steps:

a) Write atomic symbols for the central and terminal atoms.

Hydrogen atoms are always terminal Central atoms are generally those with the lowest electronegativity Carbon atoms are always central

b) Connect the central atom with each of the terminal atom by drawing a single bond.

3. For each single bond, subtract two electrons from the total number of valence electrons.
4. Using the remaining valence electrons, complete the octets of the terminal atoms first, and then complete as many as possible for the central atoms.
5. If you have used up all of the valence electrons to complete octets for all of the atoms, you are done.
6. If not, then complete the octets of all central atoms by moving lone-pairs from terminal atoms to form multiple bonds.
7. Calculate the Formal Charges on all atoms and label the non-zero formal charges in the structure:

Formal Charge on an atom = No. of valence electrons in free atom–No. of lone pair electrons –½ (No. of bonding electrons)

Formula 1.

Examples : Here we will take CO molecule as an example to explain the procedure step by step:

1. Total number of valence electrons: 4 (C atom) + 2×6 (2 O atoms) = 16

Always DOUBLE CHECK: In the correct Lewis structure, the total number of electrons involved (bonding plus non-bonding electrons) must be equal to this number, less or more are both incorrect!!

2. Write a plausible skeletal structure:

Carbon atoms are always central, so the skeletal structure is: O — C — O

3. Four electrons are used so far, and there are 16 – 4 = 12 electrons remained.
4. The remaining 12 electrons must be used to complete the octet for both terminal O atoms first, and no electrons left after that.

It is very important to keep in mind that the remaining electrons should be used to give the octet of terminal atoms first!

5. The central C atom does not get octet yet, we should do next step.
6. Moving one lone pair from each terminal O atom, the following structure is obtained.

this is the complete Lewis structure of CO.

For Lewis structure purposes, the lone-pairs can only be moved f rom terminal atoms to the central atom to form multiple bonds, not the other way around.

7. Formal charges check: all atoms have formal charges equals to 0 in this structure.

FC (C) = 4 -½×(4×2) = 0

FC (O) = 6 -4-½×(2×2) = 0

Since the two oxygen atoms have the same bonding, one calculation is enough for both oxygen atoms.

1.2.3 Guidelines about Formal Charges in Lewis Structures

The purpose of formal charges is to compare the difference between the number of valence electrons in the free atom and the number of electrons the atom “owns” when it is bonded. The smaller the difference the “happier” (more stable) the atom is. The

4 +

2

2

1.2.3 https://chem.libretexts.org/@go/page/

atom owns all of the lone pair (non-bonding) electrons and half of the bonding (shared) electrons, which is why the formula is in the way given in Formula 1.1.

Formal charges can be used as guidelines to determine the plausibility of Lewis structures by comparing the stability of non- equivalent resonance structures, which is particularly important for organic species. The rules about formal charges are:

The sum of the formal charges must equal to the total charge on the molecule or ion. Formal charges should be as small as possible (comparing the absolute value of formal charges for such purposes). “-” FC usually appears on the most electronegative atoms (with the stronger ability to pull the shared electrons; this atom is “winning” electrons in the sharing). “+” FC usually appears on least electronegative atoms (with the weaker ability to pull the shared electrons; this atom is “losing” electrons in the sharing). Structures having formal charges of the same sign on adjacent atoms are unlikely.

There is a derived way for calculating formal charge: since each bond contains 2 electrons, half of the bonding electrons simply equals to the number of bonds. So, the formal charge can also be calculated based on the derived version of the formula:

Formal Charge on an atom = No. of valence electrons in free atom–No. of lone pair electrons – No. of covalent bonds around the atom Formula 1.

Double bonds count as 2 and triple bond count as 3 in Formula 1.2. Both Formula 1.1 and 1.2 work for counting the formal charge; you can choose either one for your convenience. While almost all of the other textbooks show Formula 1.1 as the official way, Formula 1.2 is easier to use and can be regarded as the practical one based on experience.

Exercises 1.

Why is the following structure not the best way to show the Lewis structure of CO?

Answers to Practice Questions Chapter 1

1.2.4 Kekulé Structures vs Lewis Structures

The complete Lewis structure always has to include all the bonding electrons and lone pair electrons. However, organic species are usually shown as KeKulé structures (more discussion will be in Chapter 2 ) with all the lone pair electrons completely omitted (with exceptions to the lone pairs that are shown to highlight special properties). Therefore, when viewing Kekulé structures, it is very helpful to keep in mind that atoms other than C and H should have a certain number of lone pairs. Examples of Kekulé structures of some compounds are given here:

Figure 1.2f The Kekulé structures of ethanol, acetic acid, ethyl amine, and ethyl bromide To count how many lone pairs should be involved on a certain atom, apply the octet rule. All of the atoms (except H) should have 8 electrons around it, therefore, N usually has 1 lone pair, O has 2 lone pairs and halogens have 3 lone pairs.

1.2.5 Exceptions to Octet Rule in Lewis Structure

So far we have always been applying the octet rule in Lewis structures, however there are some cases in which the rule does not apply. For example, H only needs 2 electrons. Here we will see some other cases where the octet rule is compromised.

Odd number of electrons

If the total number of valence electron is an odd number, the octet rule can not be applied to all atom in the species. The examples could include NO (nitrogen monoxide or nitric oxide), NO (nitrogen dioxide) and alkyl radicals.??

NO molecule : Although NO is a diatomic molecule, it is possible to draw the Lewis structure by following the procedure. Depending on which atom is given the octet first in Step 4, you may get two possible structures. By applying the formal charge guideline, we can decide that the first structure is the better choice with zero formal charges.

2

2

1.2.5 https://chem.libretexts.org/@go/page/

Elements in Period 3 (or higher) have 3 (or more than 3) principle shells, so the d orbital is available in the valence shell. That is why they can accommodate more than 8 electrons.

Key Takeaways

For elements in 2 period, C, N, O, F and Ne, the maximum number of electrons involved in Lewis structure is eight!!!

This page titled 1.2: Lewis Structure is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by Xin Liu (Kwantlen Polytechnic University).

nd

1.3.1 https://chem.libretexts.org/@go/page/

1.3: Resonance Structures

In the case that more than one reasonable (plausible) Lewis structure can be drawn for a species, these structures are called resonance structures or resonance contributors. Resonance structures can be either equivalent or non-equivalent.

Equivalent Resonance Structures

Let’s consider the example of carbonate anion, CO

Figure 1.3a Versions of the carbonate anion Lewis structure

By following Step 6 in the Lewis structure drawing procedure , the double bond can be built between the central C and any of the terminal O’s to generate three structures, and they all look “the same”. However, they are not really identical (or same), they are just equivalent. Each structure is called a resonance structure, and they can be connected by the double-headed resonance arrow. There are total three equivalent resonance structures for CO , and the actual structure of CO is the hybrid of the three resonance contributors.

Figure 1.3b Three equivalent resonance contributors of carbonate anion The arrows used here to connect between resonance structures is the “ resonance arrow “, which has double arrow heads. Resonance structures have to be connected using resonance arrows.

Since the resonance structures are equivalent, they are all in the same level of energy and have the same stability, so they make the same contributions to the actual structure of CO. This is supported by the experimental evidence that all the carbon-oxygen bonds in CO are the same bond length, which is longer than a regular double bond but shorter than a single bond. As a result of the resonance structures, the two negative charges in CO are not localized on any oxygen atoms, but are spread evenly among all three oxygen atoms, and is called charge delocalization. Because of charge delocalization, each oxygen atom has two-thirds of a full negative charge. Charge delocalization helps to stabilize the whole species. The stability a species gains from having charge delocalization through resonance contributors is called resonance stabilization effect. The greater the number of resonance contributors, the greater the resonance stabilization effect, and the more stable the species is.

The actual structure of the carbonate anion is a combination of all the three equivalent resonance structures, that can be called a hybrid. What does the actual structure look like, and can we draw one structure on paper to show the actual structure? The actual structure can not be shown with a conventional Lewis structure, because the regular Lewis structures do not include partial charges, and there is two-thirds of a full negative charge on each oxygen atom in CO. An attempt to show the hybrid structure can be by using dashed lines to show that the bond between carbon and oxygen is somewhere between a single and double bond, and each oxygen atom has partial charges.

Figure 1.3c Dashed lines drawn on the CO3 molecule Lewis structure to show the actual structure and partial charges

3 2-:

3 2-^3 2-

3 2- 3 2- 3 2-

3 2-