Organic Chemistry Module 1 Study Guide PDF

22. 6. 19. 오후 3:31 1.1 Introduction to Organic Chemistry: Principles of Organic Chemistry with Lab-2021-Gallaher 1.1 Introduction to Organic Chemistry https://portagelearning.instructure.com/courses/1430/pages/1-dot-1-introduction-to-organic-chemistry?module_item_id=150728 1/19 22. 6. 19. 오후 3:31 1.1 Introduction to Organic Chemistry: Principles of Organic Chemistry with Lab-2021-Gallaher Introduction to Organic Chemistry Organic chemistry is a subdiscipline of the chemical sciences. The central focus of organic chemistry is the element carbon and the compounds formed by combining carbon with a relatively small number of other elements ( 10-12). The unique nature of the carbon atom allows for an amazingly diverse array of patterns of combinations between carbon and other elements. Carbon atoms are joined by forces called bonds into structures known as molecules. Molecules can contain a single or multiple carbon atoms joined together in chains or rings. The addition of hydrogen, oxygen, nitrogen, sulfur, and halogen atoms, among others, gives molecules characteristic properties and reactivity. Since molecules are formed through known principles of chemical combination, the number and type of atoms in each combination result in predictable variations of composition and geometry (shape). Currently, an organic compound is defined as a covalently bonded compound containing the element carbon, excluding carbonates and oxides (such as CaCO3 and CO2). As of today, there have been more than 20 million unique organic molecules discovered and described in the scientific literature – far greater than the comparable number of inorganic compounds. This figure speaks to the unique nature of the carbon atom and its ability to combine in multiple ways with other chemical elements. One specific reason for the large number of unique organic molecules is the propensity of the carbon atom towards catenation – the process or preference of an element to link/bond with another atom of the same element. To put it simply, carbon is an element that likes to form bonds to other carbon atoms. This results in many possible variations in the carbon skeleton or framework of the molecule, onto which other elements can be added, increasing the diversity of the compound in terms of its properties and reactivity. Organic chemistry is truly a science of patterns. Chemists have realized that the properties and reactivity of an organic molecule are controlled by its composition and shape. By studying these, chemists have developed a deeper understanding of chemical processes, such as the mechanisms behind the chemical reactions that molecules undergo. An understanding of the composition and shape of molecules has also aided chemists in the search for new compounds and materials – planning and computer modeling have allowed chemists to predict and design the properties and behavior of new molecules rather than leaving these discoveries to chance. While it is certainly true that studying organic chemistry is good preparation for chemistry and biology students and those intending to pursue professional studies like medical or veterinary school (among others), there is another ancillary benefit that comes from the study of organic chemistry – the development of critical thinking skills. Solving problems in organic chemistry often involves a way of thinking that goes well beyond routine memorization of facts or the ability to recall and rearrange equations. Organic chemistry involves understanding and correlating https://portagelearning.instructure.com/courses/1430/pages/1-dot-1-introduction-to-organic-chemistry?module_item_id=150728 2/19 22. 6. 19. 오후 3:31 1.1 Introduction to Organic Chemistry: Principles of Organic Chemistry with Lab-2021-Gallaher many different pieces of information to solve a problem. Some problems in organic chemistry can be worked forwards, others must be approached backward. Some problems in organic chemistry have several possible explanations that must be evaluated to decide which explanation best models the problem under consideration. The approach that a trained organic chemist takes to any problem is to recognize similarities between unknown things and things that are known with certainty. Oftentimes this means going back to the most fundamental chemical principles, which is where this course will begin to explore the discipline of organic chemistry. In order to fully understand the nature of an organic molecule, one must be able to describe and discuss the impact of each component of the composition of any particular molecule, from the most fundamental to the most subtle or abstract. Atomic Valence Why do atoms form bonds to other atoms? What is the driving force for atoms to join together? The answer to this question is the answer to many questions in nature – energetic stability. Many times, atoms of elements are more stable when they form bonds to other elements. To say an atom is energetically stable means that the overall energy of the atom is lower when in a bonded state than in a non-bonded state. Lower energy typically means greater stability. But why? The answer lies in the most fundamental level of atomic organization – the electron configuration of the atom, and more specifically, the atomic valence. The valence of an atom is the number of bonds an atom will form to fill its valence (outermost) shell. The atom achieves stability from a full valence shell, as there is no more “desire” for the atom to make bonds. Think of a collector with one empty space in their collection – they would likely be trying hard to secure the last piece for their collection! Knowing the valence of an atom is important because it indicates how many bonds a given atom will typically form when it combines with other elements to make a molecule. The valence for an atom can be determined by writing out the electron configuration for that atom and counting how many electrons are necessary for the atom to have a full valence shell. This is also typically referred to as a “noble gas configuration”. The elements in Group Eight of the Periodic Table of the Elements (PTOTE) are known as the noble gases due to their inert nature (non-reactive/non-bonding). For most main group elements, which are elements in the PTOTE other than the transition metals, lanthanides, and actinides, a full valence shell is composed of eight electrons. The main group elements follow the “octet rule” in trying to complete their valence shell. Hydrogen, the smallest element, has a valence shell composed of only two electrons. It is rather tedious to write out the entire electron configuration of an element to figure out how many electrons the atom needs to fill its valence shell, especially when considering elements with larger atomic numbers (and hence, more subshells in the electron configuration). Luckily, there is a method to quickly https://portagelearning.instructure.com/courses/1430/pages/1-dot-1-introduction-to-organic-chemistry?module_item_id=150728 3/19 22. 6. 19. 오후 3:31 1.1 Introduction to Organic Chemistry: Principles of Organic Chemistry with Lab-2021-Gallaher determine the valence of any element in the main group – a method that allows us to generate a visual/pictorial representation of only the valence shell of an atom – the Lewis Dot Diagram (LDD). Generating the Lewis Dot Diagram for a main group element involves the following steps: a.) Determine the Group Number of the element in question. This is the number at the top of the column in which the element is located on the PTOTE (Figure 1.1). b.) The Group Number indicates the number of valence electrons possessed by the atom in the valence shell. c.) Write the atomic symbol for the element and imagine it is contained in a four-sided box. d.) Place dots (representing electrons) around the symbol. Use only the four sides of the box. e.) Place electrons singly at first, one per side f.) Once each side has a single electron, continue by pairing electrons until the total number of electrons is placed (equal to the Group Number). https://portagelearning.instructure.com/courses/1430/pages/1-dot-1-introduction-to-organic-chemistry?module_item_id=150728 4/19 22. 6. 19. 오후 3:31 1.1 Introduction to Organic Chemistry: Principles of Organic Chemistry with Lab-2021-Gallaher https://portagelearning.instructure.com/courses/1430/pages/1-dot-1-introduction-to-organic-chemistry?module_item_id=150728 5/19 22. 6. 19. 오후 3:31 1.1 Introduction to Organic Chemistry: Principles of Organic Chemistry with Lab-2021-Gallaher Figure 1.1 The Periodic Table of the Elements. Note the Group Numbers (as Roman Numerals) at the top of the columns for the main group elements. The following examples demonstrate how to use the LDD. CARBON Carbon has the atomic symbol C and group number 4. Write the symbol C and place it in an imaginary 4-sided box (it is okay to draw a literal box too!). Now, the 4 electrons (represented by dots) are placed around the box (only along the sides). Single electrons are placed first, one per side, with no pairing initially. This results in a single, unpaired electron on each side of the box. When complete, the LDD can be used to determine two important items: 1.) The number of UNPAIRED electrons – this is the number of bonds the atom will form. By making a bond, the atom will gain an electron for its valence shell (one for each bond made). 2.) The number of PAIRS of electrons – these are called nonbonding or lone pairs. While these electrons are part of the valence shell of the atom, they do not participate in bonding. https://portagelearning.instructure.com/courses/1430/pages/1-dot-1-introduction-to-organic-chemistry?module_item_id=150728 6/19 22. 6. 19. 오후 3:31 1.1 Introduction to Organic Chemistry: Principles of Organic Chemistry with Lab-2021-Gallaher The LDD shown above for carbon illustrates that carbon typically makes four bonds in an organic molecule. This is one of the main ideas in foundational organic chemistry. Another thing the LDD of carbon reveals is that when in a molecule, typically, carbon does not have any lone pairs – all of its valence electrons are used to make bonds to other atoms. NITROGEN Nitrogen has the atomic symbol N and is found in group 5 of the PTOTE. Write the symbol N and place it in an imaginary 4-sided box (it is okay to draw a literal box still!). Now, the 5 electrons (represented by dots) are placed around the box (again, only along the sides). Single electrons are placed first one per side, with no initial pairing. Once each side has a single electron, the electrons are paired (a maximum of two per side) until the total predicted by the group number has been placed. The completed LDD for nitrogen looks like the following: The LDD for nitrogen indicates that nitrogen will typically make three bonds in an organic molecule, as it needs to obtain three more electrons via bonding to have a filled valence shell. Another thing that can be noted is that nitrogen will bear a single lone pair of electrons. While lone pairs are not used to make bonds, they are important in many other processes that can occur. Thus, recognizing the presence of lone pairs of electrons is equally as important as knowing how many bonds an atom will form. https://portagelearning.instructure.com/courses/1430/pages/1-dot-1-introduction-to-organic-chemistry?module_item_id=150728 7/19 22. 6. 19. 오후 3:31 1.1 Introduction to Organic Chemistry: Principles of Organic Chemistry with Lab-2021-Gallaher NEON One final example is neon. Neon is a member of group eight in the PTOTE, the noble gases. Recall that these elements are called noble because they do not like to make bonds to other elements. The LDD can be used to “see” why these elements are inert. The LDD for neon is shown below. What does the LDD for neon indicate about the ability of the atom to make bonds? It indicates that neon does not need to make any bonds because its valence shell is already filled (a total of 8 electrons and all electrons are paired). The neon atom has no desire to react or bond to other atoms because its valence shell is filled, and the atom is stable. The LDD of an atom is a quick, visual way to examine the valence shell of any main group element. The LDD is used to predict how many bonds an atom will form and how many lone pairs of electrons that atom will typically have when bonded together in a molecule. These are both fundamental pieces of information to understand the structures of organic molecules. Ionic, Covalent, and Polar Covalent Bonding A second important consideration to understanding molecular structure is to examine the type of bonding that occurs between the atoms in a molecule. The type of bonding that occurs between atoms in a molecule is one way of classifying a molecule and often imparts some key properties and reactivity patterns of a molecule. While the LDD indicates how many bonds an atom will form, it does not indicate the type of bonding that occurs between atoms of the elements. Ionic bonds are formed by the complete transfer of one or more valence electrons from one atom to another atom. Since electrons carry a negative electrical charge, the atom that loses electrons becomes positively charged and forms a cation. The atom that receives the electrons https://portagelearning.instructure.com/courses/1430/pages/1-dot-1-introduction-to-organic-chemistry?module_item_id=150728 8/19 22. 6. 19. 오후 3:31 1.1 Introduction to Organic Chemistry: Principles of Organic Chemistry with Lab-2021-Gallaher becomes negatively charged and forms an anion. An ion is a chemical species that possesses a non-zero electrical charge. An ionic bond is the attractive force between oppositely charged ions – an electrostatic attraction (“opposites attract”). An example of ionic bonding can be seen in the compound sodium chloride, NaCl (Figure 1.2). Figure 1.2 Formation of Sodium Cations and Chloride Anions by Electron Transfer from Sodium to Chlorine. The resulting charged species called ions are held together in a crystal lattice, which is a regular, repeating 3D arrangement of oppositely charged ions. The formation of the sodium and chloride ions shown above can be explained in terms of the valence of each atom and by the octet rule. The sodium atom has a single valence electron. To gain 7 electrons to fill its outer shell would require too much energy and would destabilize the atom. Thus, sodium can achieve a filled outer shell by giving up its single valence electron and “stepping backward” to have the same electron configuration as the neon atom – a filled valence shell/a full octet. The chlorine atom has seven valence electrons (group 7), so it only needs to gain a single electron to have a filled outer shell/completed octet. The tendency of an atom to gain or donate electrons is predictable and related to known periodic trends. Electropositive elements, such as metals on the left side of the PTOTE, tend to give up their low numbers of valence electrons and form cations, while electronegative elements, such as the non-metals that are found towards the right side of the PTOTE, tend to accept electrons and form anions. The general trend of electronegativity shows that electronegativity increases moving towards the right and going up a given group in the PTOTE. Ionic bonds form between atoms of elements with large differences in their electronegativity. Electronegativity is denoted using the Greek letter “Chi” ( ). Covalent bonds are formed by the sharing of electrons between elements with little (or no) difference in their electronegativity values. These bonds are formed by the sharing of electrons between atoms rather than the transfer of electrons from one atom to another. No ions are formed when atoms share electrons. Atoms can share one or more electrons with each other. The sharing is mutual in that each atom considers any shared electrons to be part of their valence shell, and thus, each atom can have a complete valence shell. The formation of a hydrogen https://portagelearning.instructure.com/courses/1430/pages/1-dot-1-introduction-to-organic-chemistry?module_item_id=150728 9/19 22. 6. 19. 오후 3:31 1.1 Introduction to Organic Chemistry: Principles of Organic Chemistry with Lab-2021-Gallaher molecule from two isolated hydrogen atoms is a paradigm example of a covalent bond (Figure 1.3). Figure 1.3 Formation of a Covalent Bond between Two Hydrogen Atoms. Heat is released when atoms form covalent bonds, so it is represented as a product. According to the LDD for hydrogen (Group 1), the hydrogen atom will only form one bond. In the case above, this is seen in the formation of the hydrogen molecule. In nature, hydrogen exists in its molecular (not atomic form). The reason for this is that the hydrogen atom is more stable when bonded to another atom. Elements that naturally occur in this manner are known as diatomic elements. Since the shared electrons in the molecule are attracted to the nuclei of both hydrogen atoms, the covalent bond is very strong. A balance exists between the atoms being close enough to one another to be able to share electrons and being too close to the point where the positive nuclei of the atoms begin to repel each other. The two above characteristics define two key parameters of the covalent bond: the bond strength and the bond length. The strength of a covalent bond is determined by measuring how much heat is released when a covalent bond forms (Figure 1.3). This amount of heat is different for various combinations of atoms that form a covalent bond. Heat is always released when two atoms form a bond because the atoms lose energy in the form of heat (and gain stability). To think in terms of bond strength, imagine reversing the process – how much energy would be needed to break the bond? The same amount! This amount is the bond strength. The chemical term that is used to describe the energy/strength of a covalent bond is known as the bond enthalpy. The bond length is a unit of distance that describes the perfect balance between attraction and repulsion between the atoms joined by the bond. At some equilibrium distance (which depends on which specific atoms are joining together) the atoms vibrate around each other, neither flying apart nor fusing together. This distance is the bond length. Specific values for covalent bond lengths and strengths are not typically memorized, as they can be readily looked up in reference tables. A single covalent bond is usually represented by a single line or dash and represents the pair of electrons that is shared between the two atoms joined by the bond. https://portagelearning.instructure.com/courses/1430/pages/1-dot-1-introduction-to-organic-chemistry?module_item_id=150728 10/19 22. 6. 19. 오후 3:31 1.1 Introduction to Organic Chemistry: Principles of Organic Chemistry with Lab-2021-Gallaher The hydrogen molecule also provides an example of another descriptor used to describe the covalent bond. Specifically, the H-H bond in molecular hydrogen is termed a pure (or purely) covalent bond, meaning a covalent bond that forms between atoms of identical electronegativity values. The H-H bond is classified as covalent because there is no transfer of electrons from one H to the other (no formation of ions), which results in shared electrons. Since neither atom in the bond has a high electronegativity, the discussion can be expanded to HOW the pair of electrons is shared between the atoms. Since the electronegativity difference is zero, the bonding electrons are equally shared between the two nuclei, that is, the pair resides equidistant between the nuclei of the two hydrogen atoms. Neither atom “wants” the bonding electron more than the other. A covalent bond between any two identical atoms is a purely covalent bond. This is not always the case though and often leads to a third classification of bonding. Polar covalent bonding is covalent bonding that occurs between elements with different electronegativity values but not enough of a difference that the bond becomes ionic in character. In other words, there is a sharing of electrons, but it is unequal. Electronegativity is a periodic property of the elements that is defined/determined when the element is bonded to other elements in a molecular framework. This unequal sharing results in atoms that carry a partial electrical charge, usually denoted by the symbol delta ( ). The covalent bond between the H and the Cl is what is known as polarized. The covalent bond formed between the hydrogen atom and the chlorine satisfies the valences of both atoms simultaneously by a sharing process, but due to the greater electronegativity of chlorine, the shared (bonding) pair of electrons lies closer to the chlorine than the hydrogen. Due to the greater electronegativity of chlorine, the chlorine pulls the shared pair of electrons closer to itself, resulting in a partial negative https://portagelearning.instructure.com/courses/1430/pages/1-dot-1-introduction-to-organic-chemistry?module_item_id=150728 11/19 22. 6. 19. 오후 3:31 1.1 Introduction to Organic Chemistry: Principles of Organic Chemistry with Lab-2021-Gallaher charge on the (more electronegative) chlorine and a partial positive charge on the (less electronegative) hydrogen. The polarization of the polar covalent bonds can be shown using several conventions (Figure 1.4). Bond polarization can be denoted with a dipole arrow, with the head denoting negative and the tail denoting positive (marked with a cross-hatch). The dipole arrow always points towards the more electronegative element in a polar covalent bond. Another convention uses the Greek letter delta ( ) to indicate partial electrical charge. + and - are applied to the less and more electronegative atoms (respectively) joined by the bond. The degree of polarization scales with the difference in electronegativity between the elements. Figure 1.4 Different representations of the polarization of the covalent bond in the HCl molecule. Clearly, the difference in the electronegativity ( ) values between atoms has a great deal of impact on the nature of the bond that can form between them. Large differences in electronegativity result in ionic bonds, no differences in electronegativity result in purely covalent bonds, and slight differences in electronegativity result in polar covalent bonds. There is general agreement among chemists that some benchmarks exist in classifying a bond as ionic, polar covalent, or purely covalent. The benchmarks are delineated by several key values of differences in electronegativity. The electronegativity values of some selected main group elements are shown below (Figure 1.5). https://portagelearning.instructure.com/courses/1430/pages/1-dot-1-introduction-to-organic-chemistry?module_item_id=150728 12/19 22. 6. 19. 오후 3:31 1.1 Introduction to Organic Chemistry: Principles of Organic Chemistry with Lab-2021-Gallaher Figure 1.5. Pauling Scale Electronegativity values of selected elements. The benchmarks are as follows: https://portagelearning.instructure.com/courses/1430/pages/1-dot-1-introduction-to-organic-chemistry?module_item_id=150728 13/19 22. 6. 19. 오후 3:31 1.1 Introduction to Organic Chemistry: Principles of Organic Chemistry with Lab-2021-Gallaher To determine the type of bond that would form between any two elements, one simply determines the absolute value of the difference between their electronegativity values and determines the classification based on the chart above. Note that this method does not guarantee that a bond will form but instead the type of bond that would form if one does. https://portagelearning.instructure.com/courses/1430/pages/1-dot-1-introduction-to-organic-chemistry?module_item_id=150728 14/19 22. 6. 19. 오후 3:31 1.1 Introduction to Organic Chemistry: Principles of Organic Chemistry with Lab-2021-Gallaher The type of bonding that occurs between atoms in a molecule imparts characteristic properties and reactivity to a compound. The following table summarizes some of the typical properties associated with each type of bonding (though there are always exceptions). The more polar covalent bonds a molecule has, in general, the more “mixed” the properties of the compound tend to be. Carbon and the Covalent Bonds Since the central focus of organic chemistry is on the element carbon, this section will begin to focus on the nature of the bonding between carbon and other main group elements. It was previously mentioned that carbon has four unpaired electrons in its valence shell and thus will tend to form four bonds. Being in the middle of the PTOTE, carbon is neither strongly electronegative nor electropositive. The Pauling Scale electronegativity for carbon is 2.5. In most cases in organic molecules, carbon tends to form covalent bonds with other main group elements by sharing electrons. Consider the formation of the simplest organic compound, the hydrocarbon (a molecule composed only of the elements https://portagelearning.instructure.com/courses/1430/pages/1-dot-1-introduction-to-organic-chemistry?module_item_id=150728 15/19 22. 6. 19. 오후 3:31 1.1 Introduction to Organic Chemistry: Principles of Organic Chemistry with Lab-2021-Gallaher carbon and hydrogen) methane (CH4), from a single carbon and four hydrogen atoms: The mutual sharing of electrons means that both atoms complete their valence shells. A full valence shell for hydrogen is two electrons, while for carbon, a full shell is eight electrons. The above structure for methane represents this visually. Each hydrogen has two electrons around it. The two electrons are shared between the H and the C, but the electrons are “counted” as belonging to BOTH atoms – this is the nature of a shared pair. In the same fashion, the carbon has eight electrons around it, so it has a filled octet. Since both atoms accomplish a filled valence by sharing, the compound is remarkably stable (and found in large concentrations on earth). The C-H bond is truly covalent, as the electronegativity difference between carbon and hydrogen is only 0.4. To simplify the representation of molecules whose atoms are joined by covalent bonds, the shared (bonding) pair of electrons is usually represented by a single line (sometimes called a dash) representing the two electrons in the covalent bond. Using this convention, the structure of the methane molecule can be represented as in the first image at left. As an example of some of the points from the preceding sections, as a typical covalent compound, methane is found as a gas in its natural state (at STP) and is completely insoluble in water. Recall that organic chemistry is a separate subdiscipline of the chemical sciences due in large part to the nature of the carbon atom showing a preference for forming covalent bonds to other carbon atoms (catenation). Due to this property, organic molecules can consist of multiple carbon atoms covalently bonded to one another in different patterns – linear chains, branched chains, rings, or any combination of these. An example of a “higher” hydrocarbon is the molecule ethane, a molecule composed of two carbon atoms and six hydrogens, is seen in the second (bottom) image at left. https://portagelearning.instructure.com/courses/1430/pages/1-dot-1-introduction-to-organic-chemistry?module_item_id=150728 16/19 22. 6. 19. 오후 3:31 1.1 Introduction to Organic Chemistry: Principles of Organic Chemistry with Lab-2021-Gallaher There is almost no limit to the number of carbon atoms that can be bonded together in an organic molecule. Some very large organic molecules can contain hundreds of carbon atoms covalently bonded together to form the “backbone” or “skeleton” of the molecule. The structures shown below provide examples of other hydrocarbons formed by joining different numbers of carbons atoms together in linear, branched, and cyclic (ring) patterns (Figure 1.6). Figure 1.6 Examples of Different Hydrocarbons. (a) Propane, C3H8 - a linear chain hydrocarbon. (b) 2-methylpropane, C4H10 - a branched-chain hydrocarbon. (c) Cyclobutane, C4H8 — a cyclic (ring) hydrocarbon. https://portagelearning.instructure.com/courses/1430/pages/1-dot-1-introduction-to-organic-chemistry?module_item_id=150728 17/19 22. 6. 19. 오후 3:31 1.1 Introduction to Organic Chemistry: Principles of Organic Chemistry with Lab-2021-Gallaher While each of the above examples contains differences in the number of atoms and the way the atoms are connected, the similarities of how carbon and hydrogen share electrons can be seen such that every atom in the molecule possesses a filled valence shell. Bond Order: Single and Multiple Covalent Bonds In all the previous examples, atoms have shared only a single pair of electrons with another atom to complete their valence shells. Often, atoms will share more than one pair of electrons with the same atom (with the same goal of filling the valence shell). The sharing of two (or more) pairs of electrons between two atoms results in the formation of a double (or multiple) bond between the atoms. The concept of multiple bonding is also known as bond order– the number of bonds between a pair of atoms in a molecule. In the hydrocarbon examples from the previous section, the carbon atoms in any of those molecules had a bond order of one. This is because each carbon shared only a single pair of electrons with any other given atom to which it was bonded. In organic molecules, carbon can, and often does, possess a higher bond order by making double or triple bonds to other carbon atoms. In the example shown below, a molecule consisting of two carbons demonstrates increasing bond order by forming a single, double, and triple bond between the two carbon atoms respectively (Figure 1.7). Figure 1.7 Examples of Organic Molecules that Demonstrate Increasing Bond Order of Carbon. (a) Ethane, C-C single bond (bond order of one). (b) Ethene, C-C https://portagelearning.instructure.com/courses/1430/pages/1-dot-1-introduction-to-organic-chemistry?module_item_id=150728 18/19 22. 6. 19. 오후 3:31 1.1 Introduction to Organic Chemistry: Principles of Organic Chemistry with Lab-2021-Gallaher double bond (bond order of two). (c) Ethyne, C-C triple bond (bond order of three). A multiple bond in an organic molecule is known as a site (or unit) of unsaturation. Sites of unsaturation occupy the valences that could be used to make bonds to hydrogen atoms. While these molecules are certainly different in terms of their structural formulae, in each structure, each carbon atom is completing its valence by sharing a total of eight electrons (filled octet) and each hydrogen is completing its valence by sharing a total of two electrons. This is the same pattern observed in the previous examples of hydrocarbons. What is different here is the number of pairs being shared between the two carbon atoms. Another noticeable difference (and one that is somewhat easy to rationalize) is that the examples with C-C multiple bonds have fewer hydrogen atoms. This is known as unsaturation in an organic molecule. Unsaturated compounds are molecules that possess one or more double or triple bonds (or a ring) as a part of their structure. Since the carbons are making more than one bond to another carbon, there are fewer bonds that can be made with hydrogen atoms. Remember that each line in the structural formulae representation above represents two electrons, so a double bond (two lines) represents four total shared electrons, and a triple bond (three lines) represents a total of six shared electrons. As will be discussed later on, molecules that possess multiple bonds have different (typically greater) chemical reactivity as compared to molecules with only single bonds due to the dense packing of electrons between atoms joined by multiple bonds. Atoms that only make one bond (hydrogen and the halogens) cannot participate in multiple bonding. Any of the main group elements that can form more than one bond can (and often do) form multiple bonds, which contributes to the diversity of structures in the world of organic compounds. https://portagelearning.instructure.com/courses/1430/pages/1-dot-1-introduction-to-organic-chemistry?module_item_id=150728 19/19 22. 6. 19. 오후 3:39 1.2 Constitutional Isomerism and Writing Structural Formulas: Principles of Organic Chemistry with Lab-2021-Gallaher 1.2 Constitutional Isomerism and Writing Structural Formulas https://portagelearning.instructure.com/courses/1430/pages/1-dot-2-constitutional-isomerism-and-writing-structural-formulas?module_item_id=150731 1/14 22. 6. 19. 오후 3:39 1.2 Constitutional Isomerism and Writing Structural Formulas: Principles of Organic Chemistry with Lab-2021-Gallaher Differences in Atomic Connectivity: The Principle of Constitutional Isomerism The molecular formula (MF) of a compound conveys the number and type of different atoms present in a compound. For example, the MF of ethane (C2H6) indicates that there are two carbons and six hydrogens in the molecule. While the MF is one important piece of information to understand the properties and reactivity of a compound, the structural formula, which, not only provides the same information as the MF but also shows the pattern of connectivity between the atoms in the molecule, is more important. A simple example would be to think about the compound water. The MF of water is H2O, so it could be said that water is a compound made up of two hydrogen atoms and one oxygen atom. However, without further explanation, it may (incorrectly) be assumed that the hydrogens were connected to each other followed by the oxygen (H-H-O), instead of the correct structure (H-O-H), with each of the two hydrogens connected to the central oxygen. It is important to be able to write the structural formula for an organic molecule to be able to predict its properties and reactivity. One challenge to writing the correct structural formula for a given molecule is the fact that most molecular formulas have different ways of arranging the specified number and type of atoms that still result in the valences of each atom being satisfied. Isomers are molecules that have the same number and type of atoms (same MF) but different arrangements of the atoms. More precisely, constitutional or structural isomers are defined as compounds with the same molecular formula that have different connectivity between the atoms in their structural formulae. Small changes in connectivity can have an enormous impact on the properties and reactivity of a molecule. Consider the example for MF C2H6O. There are two different ways to arrange/connect the atoms from this formula into a structural formula that satisfies the valences of all of the atoms: https://portagelearning.instructure.com/courses/1430/pages/1-dot-2-constitutional-isomerism-and-writing-structural-formulas?module_item_id=150731 2/14 22. 6. 19. 오후 3:39 1.2 Constitutional Isomerism and Writing Structural Formulas: Principles of Organic Chemistry with Lab-2021-Gallaher In each of the structures above, each carbon makes four bonds, each hydrogen makes one bond, and each oxygen makes two bonds (and has two lone pairs), as predicted by the LDD of each atom. In other words, all of the atoms have complete valences. The two molecules are quite different in terms of the connectivity of their atoms though. Ethyl alcohol features a C-C-O “core” with the hydrogen atoms surrounding it. In contrast, dimethyl ether features a C-O-C “core” with the hydrogens surrounding it. The alcohol has a hydrogen atom attached directly to the oxygen, while the ether does not (only carbon attached directly to the O). These connectivity differences may seem minor but the comparison shown below of some of the physical and chemical properties of these two compounds reveals some striking differences. https://portagelearning.instructure.com/courses/1430/pages/1-dot-2-constitutional-isomerism-and-writing-structural-formulas?module_item_id=150731 3/14 22. 6. 19. 오후 3:39 1.2 Constitutional Isomerism and Writing Structural Formulas: Principles of Organic Chemistry with Lab-2021-Gallaher The two compounds have measurable differences in virtually every property. Some differences are slight (like density), but some (like boiling point) are enormous. The type of reactions that the two compounds undergo is also vastly different. As one more example, many of us have probably ingested some ethyl alcohol in the form of beer, wine, or distilled spirits. However, hopefully, no one has ingested dimethyl ether as it is toxic! All of these differences come about as a result of the differences in connectivity between the atoms. Ethyl alcohol and dimethyl ether are constitutional isomers of each other. Writing Structural Formulae A fundamentally important skill in organic chemistry is the ability to write a correct structural formula for a compound from its molecular formula. As was discussed above, the principle of constitutional isomerism can certainly complicate this task. A perfectly correct drawing of a structural formula for a given MF that satisfies the valences of all the atoms could end up being incorrect because a constitutional isomer of the intended compound could be drawn instead. Rest assured, this course will provide plenty of practice in the ability to draw correct structural formulae. The basic recipe for the structure of an organic molecule is as follows: A backbone or skeleton of carbon atoms bonded together in various patterns including the following: Continuous (linear) chains– these have the carbon atoms bonded one after another. Branched chains– these have smaller chains of carbon (branches) that sprout from the longest continuous (parent) chain. Rings (cyclic chains)– these have the ends of a carbon chain connected to each other, forming a ring, or cyclic structure. Heteroatoms– atoms other than carbon or hydrogen bonded to the carbon skeleton at various positions. When they are present, heteroatoms in organic molecules are typically bonded together in specific patterns known as functional groups. Functional groups give a molecule known properties and reactivity and allow the molecule to be classified into a particular class or family of compounds. Hydrogen atoms that complete the valences of any carbon atoms along the various chains and branches. The following will show how to draw some structural formulae for compounds that possess the molecular formula C5H12. According to their LDD, each carbon atom prefers to make four bonds and each hydrogen makes only one bond. Neither atom will have any lone pairs. Using these https://portagelearning.instructure.com/courses/1430/pages/1-dot-2-constitutional-isomerism-and-writing-structural-formulas?module_item_id=150731 4/14 22. 6. 19. 오후 3:39 1.2 Constitutional Isomerism and Writing Structural Formulas: Principles of Organic Chemistry with Lab-2021-Gallaher basic guidelines, some structural formulae can be constructed. The simplest approach is to begin by connecting all five carbons together in a continuous chain, as shown below. In continuous chains, the carbons at the end of the chain (terminal carbons) are making only one bond, while the internal carbons are making two bonds (all to other carbons). When all of the carbons from the molecular formula have been used, the structure must then be completed by bonding the hydrogens to the carbons, assigning the correct number of hydrogens to each carbon so that every carbon completes its valence (makes a total of four bonds) (Figure 1.8). Figure 1.8 Pentane, a linear hydrocarbon molecule with molecular formula C5H12. In a linear molecule, terminal carbons will need three hydrogens each to complete their valence, while each internal carbon atom only needs two. It is always a good idea to check the final structure to be sure atoms were not added (or subtracted) in the process of creating the structure. To generate constitutional (structural) isomers, possible different patterns of connectivity that Figure 1.9 2-methylbutane, a branched would also result in each atom having a completed valence need to be considered. A different hydrocarbon molecule with molecular formula possibility would be to consider a branched carbon chain. To envision a branched chain, imagine C5H12. taking a pair of scissors and clipping the end carbon off the chain above and reattaching it to one of the internal carbons in the middle of the (now shorter) chain (see in the top image at left). https://portagelearning.instructure.com/courses/1430/pages/1-dot-2-constitutional-isomerism-and-writing-structural-formulas?module_item_id=150731 5/14 22. 6. 19. 오후 3:39 1.2 Constitutional Isomerism and Writing Structural Formulas: Principles of Organic Chemistry with Lab-2021-Gallaher When hydrogens are placed around the carbons to complete their valences, different numbers of hydrogens are needed on each carbon position to complete the valence of each. Carbon makes four bonds, and the hydrogens can be distributed appropriately, according to the number of bonds each carbon is already making in the branched chain (Figure 1.9). The molecules pentane and 2-methylbutane are constitutional isomers of each other. They have identical molecular formulae but different connectivity between the atoms, making them different compounds. As could be expected, as different compounds, they have different chemical and physical properties. When writing the structural formula for organic molecules, one must be certain that changes made to the connectivity result in the formation of different compounds. For example, removing the end carbon from the original five-carbon chain and adding it to the other end simply results in a five-carbon linear chain (the same connectivity as the original). This can be visualized by imagining drawing a pencil line through the carbons in the structures below. A continuous line could be drawn through either chain, from end to end, without lifting the pencil once. In a similar fashion, if the end carbon is removed and attached to EITHER of the internal carbons of the shortened (4-carbon) chain, the same connectivity will be obtained between the carbons: https://portagelearning.instructure.com/courses/1430/pages/1-dot-2-constitutional-isomerism-and-writing-structural-formulas?module_item_id=150731 6/14 22. 6. 19. 오후 3:39 1.2 Constitutional Isomerism and Writing Structural Formulas: Principles of Organic Chemistry with Lab-2021-Gallaher To describe the connectivity of the branched chains that result in the example above, it could be stated that “there is a four-carbon parent chain, with a one-carbon branch off of the second carbon from the end”. The same language applies to both (counting from the left and the right respectively), but this means that the connectivity is identical, and the two molecules would NOT be isomers of each other (same MF and same connectivity). For each of the branched chains above, one can visualize drawing a pencil line through a maximum of four carbons without having to lift your pencil. Is there another possible isomer for the formula C5H12? Yes, typically by shortening the longest continuous chain and adding additional branches (or increasing the length of existing branches), isomers of molecules can be generated for a given molecular formula. Another isomer of C5H12 exists as a three-carbon chain with two one-carbon branches projecting from the center (seen in the image at left). A unique feature of this isomer is that the central carbon atom completes its valence by making bonds only to carbon atoms (4 in total). It does not require any hydrogen atoms to complete its valence. The 12 remaining hydrogens in the formula are used to complete the valences of the four (terminal) carbons in the structure (Figure 1.10). https://portagelearning.instructure.com/courses/1430/pages/1-dot-2-constitutional-isomerism-and-writing-structural-formulas?module_item_id=150731 7/14 22. 6. 19. 오후 3:39 1.2 Constitutional Isomerism and Writing Structural Formulas: Principles of Organic Chemistry with Lab-2021-Gallaher Figure 1.10. 2,2-dimethylpropane a branched hydrocarbon molecule that is a constitutional isomer of MF C5H12. Are there other isomers of C5H12? The answer for this formula is no, as a total of five carbons does not allow for many unique skeletal arrangements. Ring structures are not possible for this formula because a ring introduces an additional C-C bond in the structure to close the ring – think of a snake biting its own tail. This extra bond would help to complete the valences of the carbon atoms and then all 12 hydrogens could not be attached because they would not all be needed. Remember that isomers must have the same molecular formulae as one another. Molecules cannot have different numbers of atoms and still be isomeric. Abbreviated Structural Formulae The dash structural formulae shown above are very useful because they indicate the number and type of atoms present in a molecule, in addition to showing how the atoms are connected. As has been observed, knowing the atomic connectivity is critical to understanding the https://portagelearning.instructure.com/courses/1430/pages/1-dot-2-constitutional-isomerism-and-writing-structural-formulas?module_item_id=150731 8/14 22. 6. 19. 오후 3:39 1.2 Constitutional Isomerism and Writing Structural Formulas: Principles of Organic Chemistry with Lab-2021-Gallaher properties and reactivity of the molecule. The drawbacks to these dash formulae are that they are large, occupying a lot of space on the page (especially as the molecular formulae get larger and more complex), and also, dash formulae can be tedious to draw. Chemists have developed alternate ways of representing structural formulae that communicate all the critical information (atomic connectivity) but reduce the drawbacks of writing dash formulae (size, tedious nature). There are several types of abbreviated structural formulae, and each type possesses several conventions that must be understood to represent the desired molecular structure properly and correctly. Condensed formulae show all the atoms but few (or none) of the bonds in the structure. Condensed formulae are easier and faster to draw than dash formulae but require the knowledge of valence and bonding of the atoms to fully understand the structure. An example of condensed formulae can be seen by translating the dash formulae for the C5H12 isomers into condensed representations (Figure 1.11). Figure 1.11 Condensed formula representations for the hydrocarbon isomers of molecular formula C5H12. https://portagelearning.instructure.com/courses/1430/pages/1-dot-2-constitutional-isomerism-and-writing-structural-formulas?module_item_id=150731 9/14 22. 6. 19. 오후 3:39 1.2 Constitutional Isomerism and Writing Structural Formulas: Principles of Organic Chemistry with Lab-2021-Gallaher In condensed formulae, bonds between C and H are not typically shown. The same is true with most C-C bonds as well. Sometimes these bonds are shown where branching occurs (or where there could be ambiguity). Note that in the examples above, a “CH2” (methine) or “CH3” (methyl) abbreviation is understood to represent that the hydrogens are bonded to the carbon and not each other. In other words, H-C-H is represented rather than C-H-H (in the case of the CH2). A CH3CH2 abbreviation implies that the carbons are bonded to one another, the hydrogens fill in the valences around them, and so on. The condensed formulae above can be further condensed by identifying and collapsing repetitions in the structure. Repetitions in this context are two or more of the same unit in sequence. In the condensed structure for linear pentane, there are three CH2 units between two CH3 groups. In the second example of 2-methylbutane, the carbon on which the branching occurs has two CH3 groups attached. The last example, the 2.2- dimethylpropane, essentially has four CH3 groups connected to a central carbon, so each of these formulas can be represented by further condensation (Figure 1.12). Figure 1.12 “Super-condensed” formula representations for the hydrocarbon isomers of molecular formula C5H12. https://portagelearning.instructure.com/courses/1430/pages/1-dot-2-constitutional-isomerism-and-writing-structural-formulas?module_item_id=150731 10/14 22. 6. 19. 오후 3:39 1.2 Constitutional Isomerism and Writing Structural Formulas: Principles of Organic Chemistry with Lab-2021-Gallaher These representations are sometimes referred to as “super-condensed” formulae and typically do not show any bonding at all. It is evident from the figure above that condensed formulae have the advantage of smaller size and ease of drawing, but it is critical to understand the concepts of valence and bonding to use and interpret these types of formulae correctly. Bond-line formulae show all of the bonds and none of the atoms, except for heteroatoms and hydrogens attached to heteroatoms. Bond-line formulae are by far the most common type of structural formula representation used in organic chemistry. They are the fastest and easiest way to draw molecular structures, but they require some “getting used to” due to the conventions employed. In a bond-line formula, the carbon skeleton is represented using lines representing bonds between the carbon atoms, but no carbon atoms are shown. It is understood that the carbons are represented at the end of each line or at the vertex where two lines meet. Typically, bond-line formulae are written in a “zig-zag” fashion, which makes the vertices easy to see. This zig-zag pattern also mimics the actual bond angles in molecules whose carbon atoms are all forming four single bonds. Sometimes, bond-line formulae are called “bends-and-ends” formulae because a carbon atom is understood to be everywhere the line bends or ends. Some examples of bond-line formulae can be seen in Figure 1.13, which shows the conversion of the dash formulae of the C5H12 isomers into bond-line formulae. https://portagelearning.instructure.com/courses/1430/pages/1-dot-2-constitutional-isomerism-and-writing-structural-formulas?module_item_id=150731 11/14 22. 6. 19. 오후 3:39 1.2 Constitutional Isomerism and Writing Structural Formulas: Principles of Organic Chemistry with Lab-2021-Gallaher Figure 1.13 Conversion of the dash formula representation of the C5H12 isomers to bond-line formula representation. While the placement of carbon atoms in a bond-line formula is usually easy to discern, it takes a bit more practice to determine where and how many hydrogen atoms are attached to each carbon. This becomes easy as one realizes that each carbon has as many hydrogens attached as necessary to complete its valence (carbon makes four bonds typically). Thus, for each carbon in a bond-line formula, the number of bonds carbon is making should be subtracted from FOUR to determine how many hydrogens are attached to that carbon! Bond-line formulae are excellent ways to represent molecules that possess multiple bonds and cyclic structures (or both). Multiple bonds between atoms in molecules are simply represented by drawing additional lines between the atoms. Cyclic molecules are represented by https://portagelearning.instructure.com/courses/1430/pages/1-dot-2-constitutional-isomerism-and-writing-structural-formulas?module_item_id=150731 12/14 22. 6. 19. 오후 3:39 1.2 Constitutional Isomerism and Writing Structural Formulas: Principles of Organic Chemistry with Lab-2021-Gallaher drawing the appropriate geometric shape according to the number of carbons in the molecule (Figure 1.14). Figure 1.14 Examples of Bond-Line Formulae for Unsaturated and Cyclic Compounds. The molecules are (from left to right): 2- butene (C4H8), 2-methyl-1-butene (C5H10), and cyclohexane (C6H12). When molecules contain heteroatoms, those atoms are shown as part of the structure. Hydrogen atoms, which are omitted on carbons, are always shown when those hydrogen atoms are bonded directly to a heteroatom. The following examples show the bond line formulae for some simple oxygen-containing organic compounds (Figure 1.15). https://portagelearning.instructure.com/courses/1430/pages/1-dot-2-constitutional-isomerism-and-writing-structural-formulas?module_item_id=150731 13/14 22. 6. 19. 오후 3:39 1.2 Constitutional Isomerism and Writing Structural Formulas: Principles of Organic Chemistry with Lab-2021-Gallaher Figure 1.15 Examples of Bond-Line Formulae of Oxygen-Containing Organic Compounds. The compounds are (left to right): ethyl alcohol (C2H6O), dimethyl ether (C2H6O), and 2-propanone (Acetone. C3H6O). https://portagelearning.instructure.com/courses/1430/pages/1-dot-2-constitutional-isomerism-and-writing-structural-formulas?module_item_id=150731 14/14 22. 6. 19. 오후 4:17 1.3 Formal Charge, Resonance, and Curved Arrow Notation: Principles of Organic Chemistry with Lab-2021-Gallaher 1.3 Formal Charge, Resonance, and Curved Arrow Notation https://portagelearning.instructure.com/courses/1430/pages/1-dot-3-formal-charge-resonance-and-curved-arrow-notation?module_item_id=150734 1/11 22. 6. 19. 오후 4:17 1.3 Formal Charge, Resonance, and Curved Arrow Notation: Principles of Organic Chemistry with Lab-2021-Gallaher Formal Charge The Lewis Dot Diagram (LDD) is a convenient way to quickly determine the “preferred” bonding pattern of each main group element for organic molecules, but in some molecules or during chemical reactions in which bonds to atoms are made and broken, some of the atoms in a structure may not exhibit the bonding pattern predicted by the LDD. When an atom in a molecule matches the bonding pattern predicted by the LDD (number of bonds and lone pairs), it will be electrically neutral. Atoms that do not meet the bonding pattern predicted by the LDD will possess a non-zero electrical charge, either positive (+) or negative (-) depending on the pattern of electrons around it (bonding vs. lone pairs). A charged atom within a molecule is usually a center of reactivity, and so, it is helpful to be able to determine the electric charge, if any, on atoms in a molecular structure. The formal charge of an atom is the electrical charge possessed by an atom when bonded in a molecule. An example to consider is the structure for the hydronium ion (H3O)+, the acidic species formed when a molecule of water accepts a proton from an acid donor: While it is stated that the hydronium ion (as a whole) is charged, the atom on which the charge actually resides can be determined. The language used by chemists states “to which atom is the charge localized”. In fact, the sum of the formal charge on each atom is the overall https://portagelearning.instructure.com/courses/1430/pages/1-dot-3-formal-charge-resonance-and-curved-arrow-notation?module_item_id=150734 2/11 22. 6. 19. 오후 4:17 1.3 Formal Charge, Resonance, and Curved Arrow Notation: Principles of Organic Chemistry with Lab-2021-Gallaher charge on the molecule or ion. The determination of the formal charge on an atom uses the following formula: Formal Charge = Group Number of the Element - (Dots + Dashes) (Equation 1) Where: "Dots" = non-binding electrons (may be lone pairs or single [unpaired] electrons) "Dashes" = bonding electrons (covalent bonds) For the hydronium ion, this formula can be used to quickly determine that the (+) charge resides on (is localized to) the oxygen atom itself. Each hydrogen (Group 1) possesses only one bond and no lone pairs. Therefore, using Equation 1, it can be determined that each hydrogen has a formal charge of zero (electrically neutral). This is not surprising; as the LDD of hydrogen says hydrogen prefers one bond and no lone pairs, each hydrogen in the hydronium ion matches its LDD bonding pattern. The central oxygen (Group 6) is making three bonds and possesses one lone pair. Thus, according to Equation 1, the formal charge on that oxygen is +1. The LDD for oxygen predicts oxygen prefers to https://portagelearning.instructure.com/courses/1430/pages/1-dot-3-formal-charge-resonance-and-curved-arrow-notation?module_item_id=150734 3/11 22. 6. 19. 오후 4:17 1.3 Formal Charge, Resonance, and Curved Arrow Notation: Principles of Organic Chemistry with Lab-2021-Gallaher make two bonds and have two lone pairs, so since the oxygen atom of hydronium does not match its LDD bonding pattern, it makes sense that it would bear a non-zero formal charge. The localized charge on an atom can be shown by writing any non-zero charges next to the symbol for that atom in the structure: As long as an atom in an organic molecule matches its LDD bonding pattern, it will be electrically neutral. Outside of that pattern, the atom will bear a non-zero charge. Some examples for the element carbon are shown below: https://portagelearning.instructure.com/courses/1430/pages/1-dot-3-formal-charge-resonance-and-curved-arrow-notation?module_item_id=150734 4/11 22. 6. 19. 오후 4:17 1.3 Formal Charge, Resonance, and Curved Arrow Notation: Principles of Organic Chemistry with Lab-2021-Gallaher To use Equation 1 correctly, the bonding and non-bonding electrons around each atom must be shown. This is another important reason to be able to correctly draw structural formulae for organic molecules. Electron Resonance in Organic Molecules https://portagelearning.instructure.com/courses/1430/pages/1-dot-3-formal-charge-resonance-and-curved-arrow-notation?module_item_id=150734 5/11 22. 6. 19. 오후 4:17 1.3 Formal Charge, Resonance, and Curved Arrow Notation: Principles of Organic Chemistry with Lab-2021-Gallaher With some structural formulae, there may be more than one way of arranging the electrons between the atoms in the structure that each result in all the atoms achieving a filled valence shell. An example of this can be observed by examining the structural formula for the carbonate ion (CO3)2-. Carbonate is a polyatomic ion that exists as a central carbon atom connected to three oxygen atoms. The overall charge on the carbonate ion is -2. When the structure of the carbonate ion is put together, the total number of valence electrons must be taken into account, including those that are contributed by the electrical charge. The carbon contributes four electrons (Group 4), and each oxygen atom contributes six electrons (Group 6). The electrical charge of -2 means that two additional electrons must be added – one for each unit of negative charge. This means the total number of electrons to be accommodated is 24 (4 +3(6)+2). Knowing the connectivity of the atoms and following the rules of atomic valence allows the construction of a structural formula that looks like the following: This structural formula reveals that carbon completes its valence by making a total of four bonds to oxygen (2 single and one double); one oxygen atom completes its valence by making two bonds (a double bond) and possessing two lone pairs, and the remaining oxygens make one bond and possess three lone pairs of electrons each. Applying the rules for formal charge, a formal charge of -1 is localized on each oxygen making a single bond to carbon. The carbon and the doubly bonded oxygen have a formal charge of zero because they match the bonding pattern of their LDD. Since carbonate has three oxygen atoms, it is an arbitrary choice of which oxygen should make a double bond to carbon. In fact, three “equivalent” structures can be drawn for carbonate that only differ in terms of the placement of the double bond and, consequently, which oxygen atoms then bear negative formal charges: https://portagelearning.instructure.com/courses/1430/pages/1-dot-3-formal-charge-resonance-and-curved-arrow-notation?module_item_id=150734 6/11 22. 6. 19. 오후 4:17 1.3 Formal Charge, Resonance, and Curved Arrow Notation: Principles of Organic Chemistry with Lab-2021-Gallaher In each of the above structures, there is one C=O and two C-O bonds. The structures have the same connectivity between the atoms, and they differ only in the position or distribution of the electrons among the atoms. What, then, is the relationship between the three structures? Are they constitutional isomers of each other? The answer is no. The structures above have the same number and type of atoms and also the same connectivity, which means that they are not constitutional isomers. The different representations of carbonate in the example above are known as resonance structures – two or more structural formulae of a molecule (or polyatomic ion) with identical arrangements of the atoms but different arrangements of the electrons in each. A conundrum regarding resonance structures is this: “How does the carbonate ion actually exist in nature”? Physical (experimental) measurements on the carbonate ion have revealed that only one single structure for the carbonate ion actually exists! Measurements using X-ray diffraction, a technique that uses X-ray radiation to measure the length and angle of bonds between atoms in molecules, have revealed that all three carbon-to-oxygen bonds in carbonate have the same length and strength. The carbon-to-oxygen bonds in carbonate are 1.31 Å (angstroms – a unit of length, 1 Å = 1x 10-10 m) long. This is surprising because by looking at any individual resonance structure, it would have one shorter C=O and two longer C-O. The value of 1.31 angstroms is surprising in itself, as this value is in between a pure C=O (1.20 Å) and a pure C-O (1.41 Å). An explanation that tries to combine the theory of electron resonance with the observed physical data states that, in nature, carbonate (and any structure that exhibits electron resonance) actually exists as a hybrid form or resonance hybrid, which is an average or blend of all of the individual resonance structures that can be drawn for that molecule. The data for the bond lengths in carbonate seem to reinforce this idea of a hybrid structure – the short C=O and the longer C-O get blended or averaged into three intermediate length bonds. Another impact of the hybrid structure is that the two negative formal charges that are localized onto two oxygens in any individual resonance form also get blended or https://portagelearning.instructure.com/courses/1430/pages/1-dot-3-formal-charge-resonance-and-curved-arrow-notation?module_item_id=150734 7/11 22. 6. 19. 오후 4:17 1.3 Formal Charge, Resonance, and Curved Arrow Notation: Principles of Organic Chemistry with Lab-2021-Gallaher averaged. The “math” here makes a bit more sense: two negative charges averaged over three oxygen atoms is -2/3 charge on each. Theoretical experiments using computer modeling and calculations have suggested this is also true for the carbonate hybrid. The depiction of the carbonate hybrid structure below uses dotted lines to indicate partial double bond character spread equally between the carbons and each oxygen. In addition, the overall charge on the ion of -2 is shown as being equally divided among the three oxygen atoms in the structure. The three equivalent resonance forms for carbonate make equal contributions to the overall hybrid structure. The hybrid structure for carbonate indicates that the three individual resonance structures that can be drawn by redistributing the electrons do not exist individually and also that they do not simply “interconvert” from one to the other. The hybrid structure verified by experimental data and calculations shows that carbonate exists in one true form with characteristics (bond length, bond strength, formal charges, etc.) that it “inherits” from each of its individual contributing resonance structures. An analogy would be to think about a child being an average, or blend, of the characteristics they inherit from each parent. Curved Arrow Formalism- "Pushing Electrons" Organic chemists use a convention known as curved arrow notation to show the relationship between resonance structures. This convention is sometimes referred to as “pushing electrons”, as the curved arrows are used to show how the electrons in one resonance structure are redistributed to become another and so on. An example of using curved arrows to show the relationship between the carbonate ion resonance https://portagelearning.instructure.com/courses/1430/pages/1-dot-3-formal-charge-resonance-and-curved-arrow-notation?module_item_id=150734 8/11 22. 6. 19. 오후 4:17 1.3 Formal Charge, Resonance, and Curved Arrow Notation: Principles of Organic Chemistry with Lab-2021-Gallaher forms is shown below: A curved arrow with a double-barb head is showing that a pair of electrons is being moved (redistributed). The curved arrow starts at the electron pair being moved and points to where the pair ends up. Resonance typically involves moving lone pairs and/or the pairs involved in a double or triple bond. Single bonds are not involved in resonance, as moving a single bond pair would result in a change of connectivity in the structure, which is not resonance. The double-headed straight arrow in between each pair of structures indicates that they are resonance structures of each other and contributors to the hybrid structure. Sometimes, drawing resonance structures (redistributing electrons) can result in structures that have atoms with incomplete valences. Consider the resonance structures shown below for the compound formaldehyde (CH2O): https://portagelearning.instructure.com/courses/1430/pages/1-dot-3-formal-charge-resonance-and-curved-arrow-notation?module_item_id=150734 9/11 22. 6. 19. 오후 4:17 1.3 Formal Charge, Resonance, and Curved Arrow Notation: Principles of Organic Chemistry with Lab-2021-Gallaher Structure “a” looks “better” to us than structure “b” because this structure has complete valences for all of the atoms. Each atom is meeting the bonding pattern predicted by its LDD, and thus, each atom has a zero formal charge (electrically neutral). The resonance structure “b”, formed by moving the double bond electrons and localizing them onto the oxygen initially looks “wrong”, as structure “b” then has carbon with an incomplete valence (and thus a formal charge of +1) and an oxygen with a complete valence, but a formal charge of -1. Note that the sum of the formal charges in “b” is zero, which is consistent with the overall charge on the structure. While there are distinct differences between these two structures, it is important to recognize that both are correct resonance structures. They have the same arrangement/connectivity of the atoms and the same overall number of electrons (none were added or removed). Since “a” and “b” differ only in the distribution of the electrons, they are resonance structures. Organic chemists do not use the language of “good” vs. “bad” structures but instead use terminology that describes how much of a contribution each individual structure makes to the overall hybrid structure of the compound. Since structure “a” does have many good qualities as mentioned above, this structure would be termed a major contributor to the overall hybrid structure, while “b”, which possesses many of the aforementioned problems, would be termed a minor contributor to the overall hybrid. Major contributors are more representative of the overall hybrid structure than minor contributors. The final hybrid structure will look much more like “a” than “b”. When drawing resonance structures, as long as the total number of electrons is not changed and the valence of any atom is not exceeded, then the resonance structure is a valid resonance form and how strongly it contributes to the overall hybrid structure can be judged. The following example for formaldehyde shows how a resonance structure could be drawn incorrectly: https://portagelearning.instructure.com/courses/1430/pages/1-dot-3-formal-charge-resonance-and-curved-arrow-notation?module_item_id=150734 10/11 22. 6. 19. 오후 4:17 1.3 Formal Charge, Resonance, and Curved Arrow Notation: Principles of Organic Chemistry with Lab-2021-Gallaher In this example, a lone pair from the oxygen is being moved to create a triple bond between the oxygen and the carbon. The resulting structure “b” would have the structure and formal charges as shown. At first, structure “b” seems odd just as it did in the previous example, and upon closer inspection, it is clear that this resonance move is not possible, as it would result in the carbon atom having more than four bonds. The valence of carbon has been exceeded – more than 8 electrons are around it, exceeding its octet, so this is impossible. Structure “b” cannot exist. Some atoms, notably sulfur and phosphorus, can expand their octet because they have access to empty orbitals from the d-block (and thus have a place to put extra electron pairs). Structure “b” above would be termed a non-contributor to the overall hybrid structure of formaldehyde and the hybrid would possess none of the characteristics of “b” as a result. https://portagelearning.instructure.com/courses/1430/pages/1-dot-3-formal-charge-resonance-and-curved-arrow-notation?module_item_id=150734 11/11 22. 6. 19. 오후 4:34 1.4 Atomic Hybridization; Sigma and Pi Bonding in Molecules: Principles of Organic Chemistry with Lab-2021-Gallaher 1.4 Atomic Hybridization; Sigma and Pi Bonding in Molecules https://portagelearning.instructure.com/courses/1430/pages/1-dot-4-atomic-hybridization-sigma-and-pi-bonding-in-molecules?module_item_id=150735 1/11 22. 6. 19. 오후 4:34 1.4 Atomic Hybridization; Sigma and Pi Bonding in Molecules: Principles of Organic Chemistry with Lab-2021-Gallaher Theory of Atomic Orbital Combination — Quantum Mechanics The concepts that have been discussed thus far (bonding, formal charge, resonance, etc.) have very complex origins in quantum mechanics, which is the branch of science that deals with the mathematical description of the motion and interaction of subatomic particles. To understand why atoms come together and make bonds in molecules, it is not necessary to become an expert in quantum mechanics. However, to be able to extend the material discussed thus far – specifically to learn how different patterns of connectivity lead to specific molecular geometries (shapes), how atoms actually share electrons needs to be understood. This is the basis of quantum mechanical theory as it applies to bond formation between atoms. In the quantum mechanical view of bond formation, bonds between atoms form when atoms overlap their individual atomic orbitals to create molecular orbitals. To expand on this definition a bit, an orbital is a region of space around an atom in which there is a high probability of finding an electron (> 90%). Orbitals have definitive shapes. For the main group elements involved in most organic molecules, the elements possess only s and p atomic orbitals. The s atomic orbitals are spherical, while p orbitals are dumbbell-shaped (two lobes). Within a given period on the PTOTE, only a single s orbital exists, while there are three p orbitals, each of equivalent energy but oriented in different ways in 3-D space (think of alignment along mutually perpendicular x, y, and z axes) (Figure 1.16). https://portagelearning.instructure.com/courses/1430/pages/1-dot-4-atomic-hybridization-sigma-and-pi-bonding-in-molecules?module_item_id=150735 2/11 22. 6. 19. 오후 4:34 1.4 Atomic Hybridization; Sigma and Pi Bonding in Molecules: Principles of Organic Chemistry with Lab-2021-Gallaher Figure 1.16 Shapes of the a) “S” and b) “P” Atomic Orbitals Relevant to the Main Group Elements for Organic Chemistry. The orbitals are not drawn to scale to one another. Elements of increasing period number possess s and p orbitals with increasing size, meaning a “2S” orbital is larger than a “1S” orbital, and so on. The center of each atomic orbital is called a node – a place in an orbital where there is zero probability of finding an electron. The nucleus is located at the node. A maximum of two electrons can reside in any given orbital. When the electron configuration for an atom is determined, the orbital picture of the atom is described using words. For example, the electron configuration for the carbon atom in the ground state (low- energy state) is 1s22s22p2. This reveals that the “core” of the carbon atom is the 1s shell containing two electrons [1s2], and the valence shell of the carbon atom is composed of the s and p orbitals of the second period, containing two electrons in the 2s and two electrons in the 2p orbitals https://portagelearning.instructure.com/courses/1430/pages/1-dot-4-atomic-hybridization-sigma-and-pi-bonding-in-molecules?module_item_id=150735 3/11 22. 6. 19. 오후 4:34 1.4 Atomic Hybridization; Sigma and Pi Bonding in Molecules: Principles of Organic Chemistry with Lab-2021-Gallaher [2s22p2]. This aligns with and extends what was discussed earlier – that the valence of carbon contains 4 electrons (group number 4), and the atom requires four more electrons to fill its valence shell. The electron configuration allows one to see “where” those electrons reside, in terms of the atomic orbitals of carbon. Bond Formation from the Atomic Orbital Picture From the atomic orbital perspective, an atom makes a bond to another atom by the overlapping of atomic orbitals from each atom. The overlapping of atomic orbitals creates a new space for electrons to reside. This new space is called a molecular orbital, or more specifically, a bonding molecular orbital. The simplest example of this can be seen in the formation of the hydrogen molecule (H2) from two isolated hydrogen atoms. Each hydrogen atom possesses a single, unpaired electron residing in a 1s atomic orbital. Recall from earlier discussions that hydrogen (Group 1) only makes one bond. Its LDD has a single unpaired electron, and each hydrogen will accept an additional electron (from another atom) to have a filled valence shell of two. Pictorially, the combination of the atomic orbitals can appear as follows: The 1s orbitals of the hydrogen atoms overlap in a linear, end-on fashion. Think of two objects undergoing a head-on collision. The language used to describe the resulting bonding orbital is that it has circular symmetry along the bond axis, which is the axis that runs through the two https://portagelearning.instructure.com/courses/1430/pages/1-dot-4-atomic-hybridization-sigma-and-pi-bonding-in-molecules?module_item_id=150735 4/11 22. 6. 19. 오후 4:34 1.4 Atomic Hybridization; Sigma and Pi Bonding in Molecules: Principles of Organic Chemistry with Lab-2021-Gallaher hydrogen nuclei. The bonding electrons occupy space directly between the two hydrogen nuclei. An analogy to describe bonds with circular symmetry is to think of the axle of a car. The nuclei are the wheels, while the electrons are the axle itself. End-on overlap of atomic orbitals produces sigma ( ) molecular orbitals – or more simply sigma bonds. Sigma bonds are single bonds with circular symmetry around the bond axis. End-on overlap can occur from several different types of orbital overlap (Figure 1.17). Figure 1.17 Examples of End-On Overlap of Atomic Orbitals to Produce Sigma Bonding Molecular Orbitals. a) s-s overlap, b) s-p overlap. c) p-p overlap. The bond created by each has circular symmetry around the bond axis. The bonding electrons are contained in a space directly between the nuclei of the atoms joined by the bond. Is there another way that orbitals can overlap other than end-on? The answer is yes. Since the main group elements are primarily limited to the s and p orbitals, there are two primary types of orbital overlap to focus on. The second type of orbital overlap relevant to organic chemistry is https://portagelearning.instructure.com/courses/1430/pages/1-dot-4-atomic-hybridization-sigma-and-pi-bonding-in-molecules?module_item_id=150735 5/11 22. 6. 19. 오후 4:34 1.4 Atomic Hybridization; Sigma and Pi Bonding in Molecules: Principles of Organic Chemistry with Lab-2021-Gallaher termed “sideways” overlap, and it typically is associated with p-orbitals. Figure 1.17 provides an example of end-on overlap of p orbitals to create a sigma bond. The p orbitals can overlap with each other in a sideways fashion too, as shown below. The sideways overlap of p atomic orbitals creates a different type of molecular orbital, known as a pi ( ) orbital. The major difference between sigma and pi orbitals is that pi orbitals do not have circular symmetry around the bond axis. The electrons in the bond reside above and below the plane of the nuclei connected by the bond. An analogy is to think about a hot dog in a bun. The two halves of the bun represent the two lobes of the pi molecular orbital. In future modules, it will be discussed that sideways overlap of p orbitals to produce -bonds is associated with multiple bonding between atoms. In a bond unit like a C=C, one of the bonds is a sigma bond and the second (or double) bond is a -bond. Sigma bonding is used to create the carbon skeleton or framework of most organic molecules. However, when considering the electron configuration for carbon in the ground state (1s22s22p2), there appears to be a problem. Carbon needs four electrons to complete its valence. However, there are already four electrons in the 2s and 2p orbitals, so four more electrons need to be shared with other atoms to complete the valence shell. If the electron configuration is shown pictorially, the s and p orbitals would be arranged according to the overall energy they possess, as shown below. https://portagelearning.instructure.com/courses/1430/pages/1-dot-4-atomic-hybridization-sigma-and-pi-bonding-in-molecules?module_item_id=150735 6/11 22. 6. 19. 오후 4:34 1.4 Atomic Hybridization; Sigma and Pi Bonding in Molecules: Principles of Organic Chemistry with Lab-2021-Gallaher The farther the orbital extends away from the nucleus, the greater the potential energy. Thus, the orbitals possess increasing energy with increasing period number due to overall increasing size. In addition, the spherical shape of the s orbital keeps electrons closer to the nucleus than electrons in the dumbbell-shaped p orbitals (within any given period, s orbitals are lower in energy than the corresponding p orbitals). With the electron configuration for carbon shown above, the core electrons are in the 1s orbital (recall only 2 electrons can fill any given orbital). The electrons are shown in this graphical representation as single-barbed arrows (one arrow = one electron). The 2s orbital of the valence shell also contains two electrons and the remaining two electrons are distributed into separate 2p orbitals in the valence shell following the proper rules for completing electron configurations. This leaves one of the three 2p orbitals in the valence shell empty. Since four more electrons are needed to have a full valence shell, it can be inferred that carbon will make four bonds. However, the number of bonds an atom will form can be predicted by sketching out the LDD and counting the number of unpaired electrons, which for carbon appears as follows: https://portagelearning.instructure.com/courses/1430/pages/1-dot-4-atomic-hybridization-sigma-and-pi-bonding-in-molecules?module_item_id=150735 7/11 22. 6. 19. 오후 4:34 1.4 Atomic Hybridization; Sigma and Pi Bonding in Molecules: Principles of Organic Chemistry with Lab-2021-Gallaher The LDD implies that carbon has 4 unpaired electrons to make bonds to other atoms. The electron configuration shown above suggests that carbon possesses only two unpaired electrons. While both the electron configuration and the LDD suggest that carbon will form the same number of bonds, there seems to be conflicting information about where the electrons are distributed (into what atomic orbitals) and which electrons are unpaired (to make bonds). The theories about the valence shell of carbon seem to conflict. One solution to the “conflict” is known as the theory of atomic orbital hybridization. Hybridization theory suggests that atoms do not use their native (also sometimes referred to as “raw” or “pure”) atomic orbitals in the valence shell to make bonds directly but instead blend or mix them together to make new, or hybrid, atomic orbitals and then use the hybrid orbitals to make bonds. One way to think of hybridization theory is that the atoms are “internally overlapping” and combining their atomic orbitals to make new ones before overlapping those new orbitals to make bonds to other atoms. What would a “hybridization” look like for carbon? To think of hybridization as a math equation, carbon would take the 2s orbital and the three 2p orbitals and blend them to get four new hybrid orbitals (Figure 1.18). Figure 1.18 Pictorial Depiction of the Hybridization of the 2s and 2p Orbitals of Carbon. The atom blends the single 2s orbital with the three 2p orbitals to make four new hybrid orbitals. Each hybrid orbital has characteristics of the original atomic orbitals. Again, in terms of the “math” of atomic orbital hybridization, if the four atomic orbitals are blended, four hybrid orbitals must be obtained. The shape and energy of the hybrid orbitals are a blend of the original atomic orbitals used to create them. The name for the hybrid orbitals is reflective of this too, as the hybrid orbitals are called sp3 orbitals to reflect the composition as 1 part s and 3 parts p orbital. With four hybrid sp3 orbitals of equal energy, the four valence electrons from the electron configuration can be filled and result in an electron distribution that solves https://portagelearning.instructure.com/courses/1430/pages/1-dot-4-atomic-hybridization-sigma-and-pi-bonding-in-molecules?module_item_id=150735 8/11 22. 6. 19. 오후 4:34 1.4 Atomic Hybridization; Sigma and Pi Bonding in Molecules: Principles of Organic Chemistry with Lab-2021-Gallaher the earlier conflict. Carbon uses its four hybrid sp3 orbitals, each containing a single, unpaired electron to overlap with atomic orbitals from other atoms to make sigma bonds. An example of this can be seen in the methane molecule examined earlier. Carbon uses its four hybrid orbitals to create bonds to four individual hydrogen atoms – end-on overlap of a 1s orbital from hydrogen with an sp3 orbital from carbon results in the formation of a C-H sigma bond. To make a molecule of methane, this process happens four times, resulting in four covalent bonds between carbon and hydrogen to make the hydrocarbon CH4. https://portagelearning.instructure.com/courses/1430/pages/1-dot-4-atomic-hybridization-sigma-and-pi-bonding-in-molecules?module_item_id=150735 9/11 22. 6. 19. 오후 4:34 1.4 Atomic Hybridization; Sigma and Pi Bonding in Molecules: Principles of Organic Chemistry with Lab-2021-Gallaher Another noticeable feature of the hybrid atomic orbitals of carbon is the geometry or shape. The four hybrid orbitals (each containing a single, unpaired electron) are directed into a tetrahedral arrangement around the carbon nucleus. A tetrahedron has four groups separated by angles of 109.5° in 3D space (Figure 1.19). The reason for this tetrahedral arrangement is to keep the four electrons in the orbitals as far apart from one another as possible to minimize repulsion between them (like charges repel each other). When the C-H bonds form to make methane, the resulting molecular geometry is also tetrahedral, as the bonding pairs of electrons also suffer the same repulsive forces. Any carbon (or other atom) that makes four single covalent bonds will exhibit a tetrahedral geometry around that atom. https://portagelearning.instructure.com/courses/1430/pages/1-dot-4-atomic-hybridization-sigma-and-pi-bonding-in-molecules?module_item_id=150735 10/11 22. 6. 19. 오후 4:34 1.4 Atomic Hybridization; Sigma and Pi Bonding in Molecules: Principles of Organic Chemistry with Lab-2021-Gallaher Figure 1.19 Representations of Bonding in the Methane Molecule. (a) The four hybrid sp3 orbitals of carbon in their native tetrahedral state. (b) Overlap of four H1s orbitals to make four C-H sigma bonds. (c) Methane dash structural formula showing the tetrahedral geometry of the molecule. As discussed previously, atoms can blend/average their atomic orbitals and use the new hybrid orbitals to form the bonds between them. However, producing different patterns of bond order (double, triple bonds) requires different types/degrees of atomic orbital hybridization. Another important realization is that the distribution of electrons around an atom also dictates the shape (geometry) of the groups attached to that atom, which is the topic that will be discussed next in the course. https://portagelearning.instructure.com/courses/1430/pages/1-dot-4-atomic-hybridization-sigma-and-pi-bonding-in-molecules?module_item_id=150735 11/11 22. 6. 19. 오후 4:34 1.5 Molecular Geometry: Principles of Organic Chemistry with Lab-2021-Gallaher 1.5 Molecular Geometry https://portagelearning.instructure.com/courses/1430/pages/1-dot-5-molecular-geometry?module_item_id=150738 1/5 22. 6. 19. 오후 4:34 1.5 Molecular Geometry: Principles of Organic Chemistry with Lab-2021-Gallaher VSEPR Theory and Organic Molecules As mentioned in the previous examples, the pattern of electrons around an atom controls the geometry (shape) around that atom. Furthermore, knowing the shape of a molecule can help one understand its properties and reactivity (which is the ultimate goal of an organic chemist). To take a large organic molecule with many atoms through an atomic orbital analysis would ultimately result in an understanding of the correct shape of the molecule, but it would be a very time-consuming effort. Just as the LDD is a “shortcut” for understanding the valence of an atom, chemists use a similar shortcut for describing the shape around atoms in molecules known as valence shell electron pair repulsion (VSEPR) theory. VSEPR theory is a tool for determining the shape around atoms in molecules based on an analysis of the patterns of electrons around them. The name of the theory describes why it works – electrons repel each other and want to be as far apart from each other as possible. In a molecule, the electrons are shared between atoms and so must adopt geometries (bond angles) that allow pairs to be as far apart as possible given different numbers of pairs of electrons and the type of pair (bonding or lone). Since typical organic molecules are large (containing many atoms), chemists usually do not try to describe the shape of the molecule (as a whole) but instead describe the shape around each atom or groups of atoms in the structure to get a picture of what the molecule looks like in terms of shape. To use VSEPR theory to successfully determine the shape of organic molecules, the patterns of electron distribution around atoms and their associated geometries (shapes) need to be known. For the majority of the elements present in organic molecules, the maximum number of pairs around an atom will be four (main group elements following the octet rule). The pairs are broken down into the number that are bonding and the number that are lone. This is because lone pairs take up more room and cause more repulsion than bonding pairs. Another condition for the proper use of VSEPR to predict molecular shapes is that multiple bonds (double and triple bonds) are counted as a single pair of electrons since they occupy the same region of space. https://portagelearning.instructure.com/courses/1430/pages/1-dot-5-molecular-geometry?module_item_id=150738 2/5 22. 6. 19. 오후 4:34 1.5 Molecular Geometry: Principles of Organic Chemistry with Lab-2021-Gallaher The electronic geometry, which refers to the shape assumed by the pairs of electrons themselves, can be described, but usually, chemists are more interested in the molecular geometry, which is the shape assumed by the

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