1 1.1 Structure and Bonding

Chapter 1 Learning Objectives

  1. Define organic chemistry
  2. What makes carbon special?
  3. Be able to distinguish ionic and covalent bonding.
  4. Be able to distinguish nonpolar and polar covalent bonds.
  5. Be able to draw Lewis structures of organic molecules.
  6. Be able to determine the shape of a given organic molecule.
  7. Be able to calculate the formal charge on an atom.
  8. Be able to explain the hybridization theory and bonding of carbon in organic compounds.
  9. Be able to draw expanded, condensed and skeletal structures of organic compounds.
  10. Be able to draw resonance structures.

 

1.1.1 Introduction

Welcome to the world of organic chemistry! You are most probably in this class because you need it as a prerequisite for health related programs. Perhaps, you are curious and excited to learn a new branch of chemistry. You have heard good and bad stories about taking organic chemistry. You are probably returning to school after more than a decade. No matter what your circumstance may be, the pressing question has always been: Why is study of organic chemistry so important? It is because it is the study of life itself. You will be amazed with the role played by organic chemistry in the progress of our society. This article highlights some everyday applications of organic chemistry. Listen to the TedTalk that tells you why it is a smart move to take this course.

1.1.2 How did it all begin?

For a longest time, scientists assumed that organic compounds cannot be made from non-living matter. This became popularly known as the ‘vital theory’, an unseen vital force that helped make organic compounds from living things. Friedrich Wohler, a 28 year old chemist, synthesized the first organic compound in 1828 by heating ammonium cyanate, an inorganic compound, to produce urea, an organic compound (scheme 1.1.1). This brought an end to the vital theory.

Scheme 1.1.1: Pyrolysis of Ammonium Cyanate

What is the traditional definition of organic chemistry?

Organic chemistry is the study of the synthesis, structure, reactivity and properties of a large group of chemical compounds primarily consisting of carbon. Watch this CrashCourse video to obtain additional insight into the definition of organic chemistry.

What is so special about carbon?

The answer to this question comes from carbon’s electronic structure and its consequent position in the periodic table (figure 1.1.1). As a group 4A element, carbon can share four valence electrons and form four strong covalent bonds. Furthermore, carbon atoms can bond to one another, forming long chains and rings. Carbon forms compounds with multiple bonds (double and triple bonds). Due to these characteristics of carbon, it is able to form an immense diversity of compounds, from the simple methane, with one carbon atom, to the staggeringly complex DNA, which can have more than 100 million carbons.

 

Figure 1.1.1: Position of Carbon

Here is a quick look at key terms you learned in fundamental chemistry course. These terms will be used often in organic chemistry.

Atomic number

Atomic mass

Mass number

Isotopes

Core and valence electrons

Electron configuration

 

Exercises

  1. What is the ground-state electron configuration of each of the following elements:

    (a) Oxygen (b) Nitrogen (c) Sulfur

  2. How many electrons does each of the following biological trace elements have in its outermost electron shell?

           (a) Magnesium (b) Cobalt (c) Selenium

Answers to Exercises (will be available soon)

1.1.3 Ionic bonds and covalent bonds

Elements bond with each other so that they form stable compounds. There are two fundamental kinds of chemical bonds – ionic bonds and covalent bonds. A positively charged cation formed from the element on the left side of the periodic table attracts a negatively charged anion formed from the element on the right side of the table. An example is sodium chloride, NaCl. Ionic bonds are formed from the transfer of electrons from sodium atom to chlorine atom (figure 1.1.2).

Figure 1.1.2: Formation of an Ionic Bond

Credits: Wikimedia Commons; This file is licensed under the Creative Commons Attribution-Share Alike 4.0 International license.

On the other hand, covalent bonds result from sharing of electrons. In the case of HF, both hydrogen and fluorine atoms share an electron each to form a bond between them (figure 1.1.3). Ionic bonds occur between metals and nonmetals whereas covalent bonds occur between nonmetals. Covalent bonding is the mode of bonding for organic compounds.

 

Figure 1.1.3: Formation of a Covalent Bond

Credits: Wikimedia Commons; This file is licensed under the Creative Commons Attribution-Share Alike 4.0 International license.

Watch this TedEd video that explains why atoms bond? How do they decide if they should bond ionically or covalently?

Key Takeaways: Chemical Bonding

  1. Bonding is the joining of two atoms in a stable arrangement.
  2. Through bonding, atoms attain a complete outer shell of valence electrons.
  3. Through bonding, atoms attain a stable noble gas configuration.
  4. Ionic bonds result from the transfer of electrons from one element to another.
  5. Covalent bonds result from the sharing of electrons between two nuclei.
  6. Organic compounds are formed by covalent bonding.

1.1.4 Electronegativity

Electronegativity is the property of an atom to attract the shared pair of electrons. Electronegativity values are used as a guideline to indicate whether the electrons in a bond are shared equally or unequally between two atoms. Sharing of electrons is similar to two people playing tug of war, who pulls the hardest.

Linus Pauling introduced this fundamental idea of electronegativity and formulated a scale of 0 to 4, where 4 is the value given to the most electronegative atom (figure 1.1.4).

Figure 1.1.4: Pauling’s Scale of Electronegativity

Credits: Electronegativity and Polarity. (2020, October 12). Retrieved June 28, 2021, from https://chem.libretexts.org/@go/page/210716; ” Electronegativity and Polarity” by LibreTexts is licensed under CC BY .

When electrons are equally shared, the bond is said to be nonpolar covalent. For example, in a chlorine molecule, electrons are shared equally between two chlorine atoms (figure 1.1.5).

Figure 1.1.5: Chlorine-Chlorine Non-polar Covalent Bond

On the other hand, in the H—Cl bond, the electrons are pulled away from H (2.1) toward Cl (3.2), the element of higher electronegativity (figure 1.1.6). The bond is polar, or polar covalent. The bond is said to have dipole; that is, a separation of charge. When differences in electronegativity result in unequal sharing of electrons, the bond is polar, and is said to have a “dipole”.

Figure 1.1.6: Hydrogen-Chlorine Polar Covalent Bond

Credits: Electronegativity and Polarity. (2020, October 12). Retrieved June 28, 2021, from https://chem.libretexts.org/@go/page/210716; ” Electronegativity and Polarity” by LibreTexts is licensed under CC BY .

The direction of polarity in a bond is indicated by an arrow with the head of the arrow pointing towards the more electronegative element. The tail of the arrow, is drawn towards the less electronegative element. Alternatively, a partial positive charge sign, deltaa plus, is placed on the less electronegative element and a partial negative charge, delta minus, is placed on the more electronegative element. It is important to remember that fluorine is the most electronegative element and has a value of 4.0 on the Pauling’s scale of electronegativity. Oxygen is the second most electronegative element. It is important to note that although there is an electronegativity difference between carbon and hydrogen, the difference is too small, and hence the bond is considered non-polar.

To determine the nature of the bond, one may use the following scale (figure 1.1.7 and figure 1.1.8).

Figure 1.1.7: Scale to Determine the Nature of a Chemical Bond

Figure 1.1.8: Types of Chemical Bond

 

Additional Resource:

View the following video to understand the relationship between electronegativity and polarity of molecules.

Credits: Crash Course Chemistry; Crash Course Chemistry: Crash Course is a division of Complexly and videos are free to stream for educational purposes.

The O-H bond is a common bond found in many organic molecules. It is a polar covalent bond and the polarity of -OH bonds is responsible for imparting many interesting properties to organic molecules. Learn more about the strange behavior of water due to O-H bonds from this video:

Credits: Ted-Ed; Ted-Ed is an extension of Ted and its mission is creating lessons worth sharing.

Watch these helpful free-access videos from Khan Academy on various aspects of chemical bonding.

Key Takeaways: Polarity and Electronegativity

  1. Electronegativity is the ability of an atom to attract a shared pair of electrons.
  2. Based on Pauling’s scale of electronegativity, fluorine is the most electronegative element with a value of 4.0. Oxygen is the second most electronegative element.
  3. If two atoms that share a bond do not differ in their electronegativities (<0.4), they form a nonpolar covalent bond.
  4. If two atoms that share a bond differ in their electronegativities (0.5 – 1.8) , they form a polar covalent bond.
  5. If the elements that form a bond differ in their electronegativities greater than 1.8, they form an ionic bond.
  6. Covalent bonds are formed between nonmetals while ionic bonds are formed between a metal and a nonmetal or between a cation and an anion.
  7. Predominant mode of bonding in organic compounds is covalent bonding.

Examples

Look at the following electrostatic potential map of methylamine, a substance responsible for the odor of rotting fish, and tell the direction of polarization of the C–N bond:

Answer:

 

Examples

Phosgene, Cl2CO, has a smaller dipole moment than formaldehyde, H2CO, even though it contains electronegative chlorine atoms in place of hydrogen. Explain.

Answer: There are two C-Cl dipoles that cancel the C=O dipole. Therefore, phosgene has a smaller dipole moment. However, in formaldehyde the two C-H bonds are nonpolar and the net resultant dipole moment is from C=O bond.

Examples

Fluoromethane (CH3F, μ = 1.81 D) has a smaller dipole moment than chloromethane (CH3Cl, μ =1.87 D) even though fluorine is more electronegative than chlorine. Explain.

Answer: Dipole moment is a product of charge and distance separating the charges. C-F is much shorter than C-Cl bond (due to better overlap of C-F orbitals). Due to a shorter distance, fluoromethane has a smaller dipole moment.

Examples

Methanethiol, CH3SH, has a substantial dipole moment (μ = 1.52) even though carbon and sulfur have identical electronegativities. Explain.

Answer: C-S bond is long and this increases the dipole moment.

 

1.1.5 Lewis Structures and Shapes of Molecules (VSEPR theory)

Lewis structures are a convenient way to show the distribution of electrons in a molecule. Drawing Lewis structures correctly is a necessary step to later on write correct organic structures using condensed and line drawings, the common methods by which organic structures are drawn.

The following table shows the 'normal' bonding modes of some important elements, often encountered in organic compounds. 

Element Carbon Nitrogen Oxygen Halogen
Number of Bonds 4 3 2 1
Lone Pairs 0 1 2 3

 

 

 

Exercise: Draw a valid Lewis structure for the following. Include all lone pair electrons and non-zero formal charges.

A) C2H6    B) C2H4   C)  C2H2    D) CH3COOH   E) CH3COCH3

Answers to Exercises (will be available soon)

1.1.6 VSEPR Model

Why do we need this model? The VSEPR theory helps a chemist to predict the electron geometry, electron geometrymolecular shape, bond angle and the hybridization of the central atom in a molecule. How does this knowledge help us? Watch this video to find out.

For a more detailed lecture on VSEPR theory, visit MIT OpenCourseWare for this recorded lecture.

Three common shapes are encountered in organic compounds. These are listed below:

*A lone pair is considered as one group. A double bond (and a triple bond) is considered as one group. An odd electron is not considered as a group.

Number of Groups around the Central Atom* Molecular Shape
4 tetrahedral
3 trigonal planar
2 linear

1.1.7 Calculating Formal Charge

All atoms in a molecule may not be uncharged. Charges affect the course of a chemical reaction, therefore, it is important to calculate the charge and know its location. Formal charge can help us do this. Formal charge is the charge assigned to individual atoms in a Lewis structure. Formal charge is calculated as follows.

Formal Charge of an atom = Number of valence electrons – Number of electrons it owns

Additional Resource:

A simple way of calculating formal charges is explained in this video.

Exercises: Determine the formal charge on the heteroatom in each given structure.

1.1.8 Wave Mechanical Model of the Atom: Our Current Understanding of the Structure of the Atom

Atoms are composed of a positively charged nucleus surrounded by negatively charged electrons that occupy three dimensional regions in space called orbitals. An electron has 95% chance of occupying an orbital. Where is the electron if it is not found in the orbital? Explore the fascinating world of quantum mechanics in this video.

Different orbitals have different energy levels and shapes. For example, s orbitals have a spherical shape and p orbitals are dumbbell in shape (figure 1.1.9). The 2s orbital is larger than the 1s orbital. Similarly, 3p orbitals are larger than 2p orbitals. However, px, py and pz orbitals are similar in size when they reside in the same shell. For example, 2px is of the same size as 2py and 2pz. The p orbitals residing in the same shell have the same energy. Such orbitals are known as degenerate orbitals. Do the p orbitals differ at all? Yes, they differ in their orientation in space.

Figure 1.1.9: s and p Orbitals
Credits: https://commons.wikimedia.org/wiki/File:S-p-Orbitals.svg

1.1.9 Hybridization in a Carbon Atom

The 2s and 2p orbitals in a carbon atom mix or hybridize to form hybridized atomic orbitals. The concept of hybridization helps us understand how carbon forms four equivalent tetrahedral bonds and, how carbon atoms form double and triple bonds with other carbon and heteroatoms.

1.1.9.1 sp3 Hybridization

An s orbital and three p orbitals on an atom can combine, or hybridize, to form four equivalent atomic orbitals with tetrahedral orientation. These tetrahedrally oriented orbitals are called sp3 hybrid orbitals. Note that the superscript 3 in the name sp3 tells how many of each type of atomic orbital combine to form the hybrid (figure 1.1.10).

 

Figure 1.1.10: Four sp3 hybrid orbitals, oriented toward the corners of a regular tetrahedron, are formed by the combination of an s orbital and three p orbitals (red/blue). The sp3 hybrids have two lobes and are unsymmetrical about the nucleus, giving them adirectionality and allowing them to form strong bonds to other atoms.

The asymmetry of sp3 orbitals arises because, as noted previously, the two lobes of a p orbital have different algebraic signs, + and –, in the wave function. Thus, when a p orbital hybridizes with an s orbital, the positive p lobe adds to the s orbital but the negative p lobe subtracts from the s orbital. The resultant hybrid orbital is therefore unsymmetrical about the nucleus and is strongly oriented in one direction. When each of the four identical sp3 hybrid orbitals of a carbon atom overlaps with the 1s orbital of a hydrogen atom, four identical C–H bonds are formed and methane results. Each C–H bond in methane has a strength of 439 kJ/mol (105 kcal/mol) and a length of 109 pm (figure 1.1.11). Because the four bonds have a specific geometry, we also can define a property called the bond angle. The angle formed by each H–C–H is 109.5°, the so-called tetrahedral angle. Methane thus has the structure shown in the following figure.

Figure 1.1.11: The structure of methane, showing its 109.5° bond angles and tetrahedral shape

 

We can picture the ethane molecule by imagining that the two carbon atoms bond to each other by head-on sigma (σ) overlap of an sp3 hybrid orbital from each (figure 1.1.12). The remaining three sp3 hybrid orbitals on each carbon overlap with the 1s orbitals of three hydrogens to form the six C–H bonds. The C–H bonds in ethane are similar to those in methane, although a bit weaker: 421 kJ/mol (101 kcal/mol) for ethane versus 439 kJ/mol for methane. The C–C bond is 153 pm in length and has a strength of 377 kJ/mol (90 kcal/mol). All the bond angles of ethane are near, although not exactly at, the tetrahedral value of 109.5°.

 

Figure 1.1.12: The carbon–carbon bond is formed by σ overlap of two sp3 hybrid orbitals.

Listen to this pencast on Sp3 Hybridization.

1.1.9.2 sp2 Hybridization

The 2s orbital combines with only two of the three available 2p orbitals. Three sp2 hybrid orbitals result, and one 2p orbital remains unchanged. Like sp3 hybrids, sp2 hybrid orbitals are unsymmetrical about the nucleus and are strongly oriented in a specific direction so they can form strong bonds. The three sp2 orbitals lie in a plane at angles of 120° to one another, with the remaining p orbital perpendicular to the sp2 plane (figure 1.1.13).

Figure 1.1.13: The structure of ethylene. One part of the double bond in ethylene results from σ (head-on) overlap of sp2 hybrid orbitals, and the other part results from π (sideways) overlap of unhybridized p orbitals (red/blue). The π bond has regions of electron density above and below a line drawn between nuclei.

Listen to this pencast on Sp2 Hybridization.

1.1.9.3 sp Hybridization

A carbon 2s orbital hybridizes with only a single p orbital. Two sp hybrid orbitals result, and two p orbitals remain unchanged. The two sp orbitals are oriented 180° apart on the right-left (x) axis, while the p orbitals are perpendicular on the up-down (y) axis and the in-out (z) axis.

 

Figure 1.1.14: The structure of acetylene. The two carbon atoms are joined by one sp–sp σ bond and two p–p π bonds.

When two sp-hybridized carbon atoms approach each other, sp hybrid orbitals on each carbon overlap headon to form a strong sp–sp σ bond. At the same time, the pz orbitals from each carbon form a pz–pz π bond by sideways overlap, and the py orbitals overlap similarly to form a py–py π bond. The net effect is the sharing of six electrons and formation of a carbon–carbon triple bond. Each of the two remaining sp hybrid orbitals forms a σ bond with hydrogen to complete the acetylene molecule (figure 1.1.14).

The following table summarizes the three types of hybridization that we discussed in detail.

To determine the hybridization of an atom in a molecule, we count the number of groups around the atom. The number of groups (atoms and non-bonded electron pairs) corresponds to the number of atomic orbitals that must be hybridized to form the hybrid orbitals.

Number of Groups around the Central Atom Hybridization Number & types of orbitals Bond Angle Molecular Shape
4 sp3 four sp3 hybrid orbitals 109.5° tetrahedral
3 sp2 three sp2 hybrid orbitals + one p orbital 120° trigonal planar
2 sp two sp hybrid orbitals + two p orbitals 180° linear

 

 

 

 

As the number of electrons between two nuclei increases, bonds become shorter and stronger. Triple bonds are shorter and stronger than double bonds, which are shorter and stronger than single bonds.

Type of Bond Bond Length A° Bond Dissociation Energy (Kj/mol)
C-C (single bond) 1.543 346
C=C (double bond) 1.483 602
C≡C (triple bond) 1.378 835
Source: Data from J. E. Huheey, E. A. Keiter, and R. L. Keiter, Inorganic Chemistry, 4th ed. (1993); John D. Robert and Marjorie C. Caserio (1977) Basic Principles of Organic Chemistry, second edition. W. A. Benjamin, Inc. , Menlo Park, CA. ISBN 0-8053-8329-8. This content is copyrighted under the following conditions, “You are granted permission for individual, educational, research and non-commercial reproduction, distribution, display and performance of this work in any format.”

Examples

Pyridoxal phosphate, a close relative of vitamin B6, is involved in a large number of metabolic reactions. What is the hybridization and the bond angle for each nonterminal atom?

Answer:

 

Key Takeaways: Hybridization Theory

  1. Hybridization theory explain bonding of carbon in organic molecules.
  2. There are three types of hybridizations for carbon: sp3, sp2 and sp.
  3. sp3 hybridization explains bonding of carbon in saturated compounds; sp2 hybridization explains bonding of carbon in unsaturated double bonded compounds while sp hybridization explains bonding of carbon in unsaturated triple bonded compounds.
  4. Hybridization explains both the structure and reactivity of organic compounds.

Exercise: Check your knowledge in hybridization theory

1.1.10 Drawing Three Dimensional Structures

In a three dimensional structure of a tetrahedral organic compound, there are three types of lines – a solid line, a solid wedge and a broken wedge. A solid line is used to indicate bond in the plane. A wedge is used to indicate a bond in front of the plane. A dashed line is used to indicate a bond behind the plane (figure 1.1.15). The molecule can be turned in many different ways, generating many equivalent representations.

Figure 1.1.15: A representation of van’t Hoff’s tetrahedral carbon atom. The solid lines represent bonds in the plane of the paper, the heavy wedged line represents a bond coming out of the plane of the page toward the viewer, and the dashed line represents a bond going back behind the plane of the page away from the viewer.

We do not need to draw molecules in three dimensional manner all the time. We have choices depending on our needs.

There are three other types of structural formulas used in organic chemistry.

  1. Expanded Formulas
  2. Condensed Formulas
  3. Line-Angle Formulas

Expanded drawings are very useful for a beginning organic chemistry student. In this structure, all bond connections are clearly shown. Lone pairs are not drawn. Although cumbersome and time-consuming, these are useful structural drawings to represent isomers.

In a condensed structure, all bonds are drawn, but single bond lines are omitted. Atoms are usually drawn next to the atoms to which they are bonded. Parentheses are used around similar groups bonded to the same atom. Lone pairs are not drawn. Sometimes, multiple bonds are also not drawn.

Challenges: Interpreting parentheses, double bonds, triple bonds, branches and how these are all attached.

Line-Angle Formulas (sometimes referred to as skeletal drawings): Skeletal structures are commonly used to represent organic compounds, in particular for cyclic compounds, because of their ease of drawing. This drawing assumes that there is a carbon atom at the junction of any two lines or at the end of any line. Hydrogen atoms are always assumed and therefore omitted in the structure. The number of missing hydrogen atoms can be determined by the tetravalency of carbon. Heteroatoms and hydrogen atoms directly bonded to them are drawn and are never assumed. This by far, is the most popular form of drawing organic structures, in particular, for cyclic compounds. Here is a summary of the main points:

  1. Each vertex represents a carbon atom
  2. C-H bonds are omitted. The number of hydrogen atoms is determined by subtracting the drawn bonds from four.
  3. Heteroatoms must be specified, and O-H and N-H bonds are not omitted. These must be clearly drawn.

Examples

Quetiapine, marketed as Seroquel, is a heavily prescribed antipsychotic drug used in the treatment of schizophrenia and bipolar disorder. Convert the following representation into a skeletal structure, and give the molecular formula of quetiapine.

Answer:

Molecular Formula: C21H25N3O2S

 

Exercises: Drawing Condensed Structures

1.1.11 Isomers

Compounds that have the same molecular formula but different arrangements are known as isomers. There are many types of isomerism. Structural isomers are those that differ in the arrangement of atoms.

Examples: Structural Isomers

Draw the indicated number of structural isomers for each of the following molecular structures.

Examples

Oxaloacetic acid, an important intermediate in food metabolism, has the formula C4H4O5. Propose two possible structures.

Answer:

Examples

The following model is a representation of citric acid, the key substance in the so-called citric acid cycle, by which food molecules are metabolized in the body. Only the connections between atoms are shown; multiple bonds are not indicated. Complete the structure by indicating the positions of multiple bonds and lone-pair electrons (black = C, red = O, gray = H).

Answer:

 

 

1.1.12 Theory of Resonance

Resonance is a concept that describes the delocalization of electrons within molecules or ions. Resonance structures are used when a single Lewis structure cannot fully describe the bonding that occurs within a molecule or ion. It involves constructing multiple Lewis structures that, when combined, represent the full electronic structure of the molecule or ion – the Resonance Hybrid. In general, molecules or ions with multiple resonance structures will be more stable than one with fewer. Some resonance structures contribute more to the stability of the molecule or ion than others – formal charges aid in making this determination.  Resonance stabilization is a major determining factor in the outcome of many reactions.

Arrows are used for drawing resonance structures with ease. It is important to remember that the resonance hybrid is the most accurate representation and structures do not flip back and forth between individual resonance structures. Watch this video for an introduction to resonance structures. This sets the stage for drawing curved arrows and generating resonance structures from a given structure.

There are several rules that must be followed when drawing curved arrows for resonance structures:

Rules for Drawing Resonance Structures: 

  1. Only electrons move!! Nuclei (atoms) and single bonds do not move.
  2. Second-row elements do not exceed octet electrons.
  3. Resonance involves electrons in lone pairs and π bonds.
  4. All resonance structures must be valid Lewis structures (ions allowed).
  5. Track all lone pairs and formal charges.
  6. Arrow pushing are used to interconvert resonance structures.
  7. Double headed arrows are used to separate resonance structures.
  8. Lower energy resonance structures contribute more to the resonance hybrid.

Recognizing common patterns greatly help with drawing correct resonance structures. The following are five patterns commonly encountered with organic structures (figure 1.1.16).

Figure 1.1.16: Common Resonance Patterns

Credits: https://www.oercommons.org/courses/homeunc-system-organic-chemistry-digital-course

Some useful resources…..

This document illustrates the correct method of drawing resonance structures.

This video explains the commonly observed patterns in resonance. 

Key Takeaways: Resonance Theory

  1. Resonance structures are multiple correct Lewis structures drawn for a particular molecule or ion.
  2. All rules for Lewis structures must be followed while drawing resonance structures. Formal charges must be correctly placed when it is other than 0.
  3. Move only pi electrons and lone pairs.
  4. Always move electrons from electron rich site to electron poor site.
  5. Never disobey the law of conservation of matter – do not create or destroy electrons. Only move them!

Exercises: Draw resonance structures for each given Lewis structure.

Among the various resonance structures, the most contributing or significant resonance structure can be determined by applying the following rules:

Rule 1: The most significant resonance contributor has the least number of atoms with formal charges.

Rule 2: The most significant resonance contributor has the greatest number of full octets (or if applicable, expanded octets).

Rule 3: If formal charges cannot be avoided, the most significant resonance contributor has the negative formal charges on the most electronegative atoms, and the positive formal charges on the least electronegative atoms.

Rule 4: The most significant resonance contributor has the greatest number of covalent bonds.

Rule 5: If a pi bond is present, the most significant resonance contributor has this pi bond between atoms of the same row of the periodic table (usually carbon pi bonded to boron, carbon, nitrogen, oxygen, or fluorine).

Rule 6: Aromatic resonance contributors are more significant than resonance contributors that are not aromatic.

 

Examples

Identify the most and the least contributing resonance structure and state your reason.

Answer:

 

Key Takeaways for Chapter 1

  1. The study of carbon compounds is known as organic chemistry.
  2. Organic compounds are formed by covalent bonding.
  3. Difference in electronegativity between atoms in shared bonds leads to polar covalent bonds.
  4. VSEPR theory predicts the shapes of molecules. The three main shapes encountered in organic compounds are tetrahedral, planar and linear.
  5. Hybridization theory explain bonding modes of carbon in single, double and triple bonded compounds.
  6. Condensed and skeletal structures are used to represent organic compounds.

 

 

definition

License

Icon for the Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License

Organic Chemistry: Fundamental Principles, Mechanisms, Synthesis and Applications Copyright © by Sujatha Krishnaswamy is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License, except where otherwise noted.

Share This Book