# Chapter 4 Lecture Notes

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1 Chapter 4 Lecture Notes Chapter 4 Educational Goals 1. Given the formula of a molecule, the student will be able to draw the line-bond (Lewis) structure. 2. Understand and construct condensed structural formulas and skeletal structures given the line-bond (Lewis) structures. 3. Given the line bond structure, determine the formal charge for O & N. 4. Define electronegativity and explain its relationship to polar covalent bonds. Give a simple rule that can be used to predict whether or not a covalent bond is polar. 5. List the five basic shapes about an atom in a compound and describe the rules used to predict the molecular shape. Explain how shape plays a role in determining overall polarity. 6. Describe the noncovalent interactions that attract one compound to another. 7. Describe the four families of hydrocarbons. 8. Explain the difference between constitutional isomers, conformations, and the stereoisomers known as geometric isomers. Give examples of two different classes of hydrocarbons that can exist as geometric isomers. 9. Define the term functional group and describe the features that distinguish hydrocarbons, alcohols, carboxylic acids, and esters from one another. Line Bond Structures We draw line bond structures to see how are bonded together in molecules. Line bond structures can be very simple as in the case of H 2, Or very complex as in the case of large molecules such as DNA. Drawing Line Bond Structures Step 1: Count the total number of from all the atoms in the molecule. Example: H 2 O 2 H atoms 2 x 1e - = 2e - 1 O atom 1 x 6e - = 6e - Total number of valence e- = 8e- The line bond structure of H 2 O will have 8 electrons 1

2 Step 2: Draw the Skeleton Structure Attach the atoms together with in the most symmetric way possible. Example: H 2 O Which of the skeleton structures is the most symmetric? Step 3: Subtract the number of electrons used to make the skeleton structure from the total number of valence electrons. How many electrons are in the skeleton structure? Step 4: Add the remaining electrons as as evenly as possible on all atoms except hydrogen. Step 5: Check for. Are there 8 electrons around all atoms (except hydrogen)? YES! If NO, use to make double or triple bonds. Step 6 (if needed): Use lone pairs to make or bonds until the octet rule is satisfied for all atoms in the molecule. 2

3 Let s do a couple of examples to see how that works! We will do O 2, N 2, and NH 3 Example O 2 Step 1: Count the total number of valence electrons from all the atoms in the molecule. How many valence electrons in O 2? Calculate valence electrons in O 2 below: Step 2: Draw the Skeleton Structure Attach the atoms together with single bonds in the most symmetric way possible. Step 3: Subtract the number of electrons used to make the skeleton structure from the total number of valence electrons. Total # of electrons in structure (from step 1 above) # of electrons used in skeleton - (from step 2 above) Remaining # electrons to be added= Step 4: Add the remaining electrons as lone pairs as evenly as possible on all atoms except hydrogen. Step 5: Check for Octets Are there 8 electrons around all atoms (except hydrogen)? If NO, use lone pairs to make double or triple bonds. Example N 2 3

4 Step 1: Count the total number of valence electrons from all the atoms in the molecule. How many valence electrons in N 2? Calculate valence electrons in N 2 below: Step 2: Draw the Skeleton Structure Attach the atoms together with single bonds in the most symmetric way possible. Step 3: Subtract the number of electrons used to make the skeleton structure from the total number of valence electrons. Total # of electrons in structure (from step 1 above) # of electrons used in skeleton - (from step 2 above) Remaining # electrons to be added= Step 4: Add the remaining electrons as lone pairs as evenly as possible on all atoms except hydrogen. Step 5: Check for Octets Are there 8 electrons around all atoms (except hydrogen)? If NO, use lone pairs to make double or triple bonds. 4

5 Draw the line bond structure for ammonia: NH 3 Step 1: Count the total number of Step 2: Draw the Skeleton Structure valence electrons from all the atoms in -Attach the atoms together with single the molecule. bonds in the most symmetric way possible. Chemical formula of molecule Central atom is: Atom # of atoms # of valence electrons x = x = x = totals Draw skeleton: Total # of electrons in line bond structure = Step 3: Subtract the number of electrons used to make the skeleton structure from the total number of valence electrons. Total # of electrons in structure (from step 1 above) # of electrons used in skeleton - (from step 2 above) Remaining # electrons to be added= # of electrons used in skeleton = (multiply # of bonds in skeleton by 2) Step 4: Add the remaining electrons as lone pairs as evenly as possible on all atoms except hydrogen. First: Re-draw skeleton here: Step 5: Check for Octets Check the structure in step 4 for octets (4 pairs) of electrons around each atom. Add remaining electrons to the skeleton as lone pairs If octet rule is satisfied, you are done. If octet rule is not satisfied, change lone pairs into double or triple bonds until octet rule is satisfied for all atoms. 5

6 Structural Formulas and Formal Charges are uncharged groups of nonmetal atoms connected to one another by covalent bonds. The structure of a molecule can be represented by an electron dot structure or a line-bond (Lewis) structure ( ). The molecular formula of isopropyl alcohol is C 3 H 8 O. The structural formula of isopropyl alcohol is: Which do you think is more informative? Structural Formulas Three completely different molecules have the same formula of isopropyl alcohol: C 3 H 8 O Structural formulas, like the line-bond formulas shown above, indicate the relative of each atom within molecules. Formal Charge In addition to describing the structure of molecules, structural formulas also provide information about where the is sometimes located in molecules and polyatomic ions. 6

7 Example: The nitrate ion (NO 3 - ), consists of a nitrogen atom covalently bonded to three oxygen atoms. The charge of the ion is shown as a superscript to the right. The charge on each atom is shown. Polyatomic Ion Examples Polyatomic ions are charged groups of nonmetal atoms that are connected by covalent bonds. (a) The net charge of the ion is shown as a superscript to the right. (b) The formal charge on each atom is shown. What do I need to know about formal charges in this class? We will keep track of formal charges on two atoms: and If oxygen has just single bond going to it, it has a formal charge of (1-). In other cases, it will have no formal charge. If nitrogen has a on it, then it will have no formal charge. In other cases, it will have a (1+) formal charge. 7

8 Group Work The line-bond structure glycine, an amino acid present in many proteins, is shown below. Assign formal charges to the nitrogen and each oxygen atom in this compound. Condensed Structural Formulas A condensed structural formula describes the attachment of atoms to one another, without showing all of the. For example, a carbon with 3 attached hydrogen atoms can be written CH 3 In condensed structures, bonds are not shown for carbon atoms that are in a sequence, or hydrogen-carbon and hydrogen-oxygen bonds. All other bonds are drawn. 8

9 Skeletal Structures In skeletal structures, covalent bonds are represented by lines, carbon atoms are shown, and hydrogen atoms are drawn only when attached to atoms other than. To read a skeletal structure, you assume that the carbon atom appears where (bonds) meet and at the of each line. Example: Draw condensed and skeletal structures of diethyl ether, a compound once used as a general anesthetic. 9

10 Group Work Research suggests that drinking green tea may help boost the immune system. Ethylamine (below), produced when one of the compounds in green tea is broken down in the liver, may be responsible for this immune response. Draw condensed and skeletal structures for this molecule. Draw condensed structure here Draw skeletal structure here SEE Chapter 4 online homework: Video: Guide to Drawing molecules - and- Practice Drawing Molecules 10

11 Polar Covalent Bonds, Shape, and Polarity Molecular Shape Definition of Molecular Shape: the three-dimensional arrangement of a molecule s in space. Molecular Shape can be determined from the line bond structure using Valence Shell Electron Pair Repulsion Theory (VSEPR). Predicting Molecular Shape VSEPR Theory The shape around the central atom(s) can be predicted by assuming that the areas of electrons (often called electron groups - E.G.) on the central atom will each other. Each counts as 1 electron group (EG) single, double or triple count as 1 EG Each counts as 1 electron group (EG) Even though lone pairs are not attached to other atoms, they do occupy space around the central atom. 11

12 Predicting Molecular Shape Step 1. Get Electron Arrangement What happens if 2 negatively charged balloons are placed into a hollow, clear plastic sphere? Electrostatic will cause the balloons to move as far apart from each other as possible. The same thing happens with. Draw the line bond structure for CO 2 in the box below: The central atom goes in the of the sphere. How many electron groups are around the central atom? There are electron groups around the central atom! Each bonded atom counts as electron group. The electron groups are placed as apart from each other as possible! 12

13 Example: 3 Electron Groups What happens if 3 negatively charged balloons are placed into a hollow, clear plastic sphere? Electrostatic repulsion will cause the balloons to move as far apart from each other as possible. Draw the line bond structure for SO 3 in the box below: The central atom goes in the middle of the sphere. How many electron groups are around the central atom? There are electron groups around the central atom! Each bonded atom counts as one electron group. The electron groups are placed as far apart from each other as possible! All electron groups are on the same and are at angles. Let s rotate the ball to see the bond angles 13

14 Example: 4 Electron Groups What happens if 4 negatively charged balloons are placed into a hollow, clear plastic sphere? Electrostatic repulsion will cause the balloons to move as far apart from each other as possible. Draw the line bond structure for a methane molecule (CH 4 ) in the box below: The central atom goes in the middle of the sphere. How many electron groups are around the central atom? There are electron groups around the central atom! The electron groups are placed as far apart from each other as possible! This arrangement is a 4-sided, 3 dimensional structure with electron groups at angles. Another way to visualize the tetrahedral shape: 14

15 Electron Arrangement Review 2 Electron Groups: 3 Electron Groups: 180 o 4 Electron Groups: EG 110 o Predicting Molecular Geometry Step 1. Get Electron Arrangement no problem now! Step 2. Determine Molecular Shape based on the arrangement of bonded. Treat lone pairs as!!! Let s do an example: Draw the line bond structure of CO 2 How many electron groups are around the central atom? What is the bond angle in a 2-electron group arrangement? Choices: a) 180 o, b) 120 o, c) 110 o 15

16 Now build a model of CO 2. Use the black sphere for the central Carbon atom. Use two metal springs for each double bond. Use the red spheres for the oxygens. This molecular shape is called. We write the general notation for this shape as. A = the central atom B = bonded atoms E = a lone pair (none seen in this example) 16

17 The molecular shape of a molecule is always! Fill in the box for 2EG with 1 lone pair on the Molecular Shape Table on the previous page. Let s do an example: Draw the line bond structure of SO 3 How many electron groups are around the central atom? What is the bond angle in a 3-electron group arrangement? Choices: a) 180 o, b) 120 o, c) 110 o Now build a model of SO 3 Use the black sphere for the central sulfur atom. Use wood sticks for single bonds and two metal springs for double bonds. Attach the red spheres with wood sticks for the oxygens. This molecular shape is called. We write the general notation for this shape as. Draw this shape and geometry name in the Molecular Shape Table on page 16. Another example: Draw the line bond structure of ozone gas O 3 Now build a model of O 3 Use the black spheres for the oxygen atoms. Use two metal springs for the double bond. Use wood sticks for the single bond. How many electron groups are around the central atom? What is the bond angle in a 3-electron group arrangement? Choices: a) 180 o, b) 120 o, c) 110 o 17

18 When determining molecular shape, we treat the lone pairs as if they were. This molecular shape is called. We write the general notation for this shape as. Draw this shape and geometry name in the Molecular Shape Table on page 15. Another example: Draw the line bond structure of methane gas CH 4 How many electron groups are around the central atom? What is the bond angle in a 4-electron group arrangement? Choices: a) 180 o, b) 120 o, c) 110 o Now build a model of CH 4 Use a black sphere for the central carbon atom. Use wood sticks for single bonds. Use yellow spheres for the hydrogens. This molecular shape is called. We write the general notation for this shape as Draw this shape and geometry name in the Molecular Shape Table on page 16. Example: Draw the line bond structure of ammonia NH 3 How many electron groups are around the central atom? What is the bond angle in a 4-electron group arrangement? Choices: a) 180 o, b) 120 o, c) 110 o Now build a model of NH 3 Use a black sphere for the central nitrogen atom. Use wood sticks for single bonds. Use yellow spheres for the hydrogens. When determining molecular geometry, we treat the lone pair as if it was invisible! 18

19 This molecular shape is called. We write the general notation for this shape as. Draw this shape and geometry name in the Molecular Shape Table on page 16. Example: Draw the line bond structure of a water molecule H 2 O Now build a model of H 2 O How many electron groups are around the central atom? Use a black sphere for the central oxygen atom. Use wood sticks for single bonds. Use yellow spheres for the hydrogens What is the bond angle in a 4-electron group arrangement? Choices: a) 180 o, b) 120 o, c) 110 o When determining molecular shape, we treat the lone pairs as if they were invisible! This molecular shape is called. We write the general notation for this shape as. Draw this shape and geometry name in the Molecular Shape Table on page

20 Polarity of Molecules Why is a molecule s polarity important? Polarity is responsible for many physical and chemical properties of substances: Freezing/melting point Boiling/condensation point How molecules interact with other molecules (including solvents and surfaces) How large molecules such as DNA and proteins fold-up on themselves How antibodies and enzymes bind with their target molecules We must know 2 things to determine if a molecule is polar: 1) No problem now!!! 2) Are there polar bonds in the molecule? 20

21 Bond Polarity Unequal sharing of electrons in a covalent bond results in bonds. Let s consider 2 simple molecules: H 2 and HCl The 2 electrons that are shared in the H 2 molecule are shared between the two hydrogen atoms. Each hydrogen atom has an equal attraction for the electrons in the bond. The 2 electrons that are shared in the HCl molecule are shared between the hydrogen and the chlorine atom. The electrons have a stronger attraction to the chlorine atom than to the hydrogen. Covalent bonding between atoms results in unequal sharing of the electrons. The result is. The electrons spend more near the Cl atom. One end of the bond has larger electron than the other. The end with the larger electron density gets a charge. The end that is electron deficient gets a partial charge. An is created by equal but opposite charges that are separated by short distances. We use an with a cross to represent an electric dipole. The arrow always points in the direction. Electronegativity The ability for atoms to attract the electrons in a covalent bond is called. Hydrogen has an electronegativity of 2.1 Chlorine has an electronegativity of 3.0 The the difference in electronegativity between the two bonded atoms, the the dipole. 21

22 Periodic Trend in Electronegativity You try it! Which molecule has a larger dipole? HF or HBr Calculate the differences in the electronegativities (E.N.) for each of the molecules. E.N. for H= E.N. for F = E.N. for H= E.N. for Br = The larger the difference in electronegativity between the two bonded atoms, the stronger the dipoletherefore HF has the largest dipole! Polarity of Diatomic Molecules Summary: That is easy for diatomic molecules, but what about the polarity of molecules with more than two atoms? 22

23 Polarity of Small Molecules We must look at what happens when we up all the dipoles. Example: CO 2 There are 2 polar bonds in an CO 2 molecule. This results in two.. The oxygens have a negative charge. The carbon has a partial positive charge. Since the molecule is symmetric, the dipoles cancel each other and the molecule is. Symmetric small molecules with are non-polar molecules. Let s look at what happens when we add up all the dipoles in another molecule: Example: H 2 O There are 2 polar bonds in an H 2 O molecule. This results in two dipoles. The oxygen has a partial negative charge. The hydrogens have a partial positive charge. Since the molecule is not symmetric, the dipoles do not cancel each other and the molecule is. Since dipoles in non-symmetric molecules do not cancel each other out, they add up to yield a molecule with a dipole. Non-symmetric small molecules with polar bonds are polar molecules. Polarity of Small Molecules: General Rule: Symmetric molecules have the central atom surrounded by electron groups. Example: 23

24 Polarity of Small Molecules Summary Chart: Polarity of Large Molecules In this course, we will be working with large molecules (organic and biochemistry). Large molecules will be considered polar (or have polar regions) if they have bonds. o The important highly-polar covalent bonds, especially in organic and biochemistry, are those in which either hydrogen or carbon atoms are covalently attached to nitrogen, oxygen, fluorine, or chlorine atoms. Understanding check The compound on the right has been used to kill any insect larvae present in cereal and dried fruit. 1) What size class would the molecule below be categorized as? a) Diatomic molecule b) Small molecule c) Large molecule 2) This molecule is. a) polar b) non polar 24

25 Polarity of Large Molecules Look for highly polar bonds Example: Cyclohexane vs. Cyclohexanol Dipoles are attracted to other dipoles by electrostatic forces! Polar molecules have strong attractions to other polar molecules because the positive end of a dipole is attracted to the negative end of another dipole. Next, will see how the polarity of molecules influences properties such as melting and boiling points of compounds. Noncovalent Interactions/Intermolecular Forces Why are some molecular compounds solid while others are gaseous and others are liquid at room temperature? Lets take a magnified look at these three phases of matter on the molecular scale! Representations of the gas, liquid, and solid states. Adding to liquids will overcome the forces holding the molecules together boiling. Adding energy to solids will overcome the forces holding the molecules together. 25

26 What forces hold molecules together to make liquids and solids? The attractive forces that hold molecules together are called forces. 3 Types of Intermolecular Forces 1) Dipole-Dipole 2) Hydrogen Bonding 3) London Forces (Induced Dipole Forces) 1) Dipole-Dipole Forces Dipoles are attracted to other dipoles by electrostatic forces! Polar molecules have strong attractions to other polar molecules because the positive end of a dipole is attracted to the negative end of another dipole. The forces of attraction between polar molecules are known as forces. 2) Hydrogen Bonding A particularly strong type of intermolecular force is called. When hydrogen is covalently bonded to a very electronegative atom, it is strongly attracted to a on another electronegative atom. Hydrogen bonding only occurs between a hydrogen bonded to N, O, or F, and a lone pair on a N, O, or F!!! 26

27 3) London Dispersion Forces If we cool a non-polar molecule such as N 2 or O 2, it will turn liquid, then solid. What forces hold these molecules together? London dispersion forces. Also known as dipole force. Caused by electrons on one molecule the electron cloud on another. All molecules have London dispersion forces (not just non-polar molecules) London dispersion forces are weaker than hydrogen bonding forces. molecules have London dispersion forces. Intermolecular Forces Conclusion Intermolecular Forces Three types: Dipole-dipole Hydrogen bonding London Dispersion Forces Much weaker than ionic or covalent bonds Larger attractive forces between molecules in pure substance means: higher point higher point 27

28 Like Like Polar molecules dissolve in polar solvents Example: Water, alcohol Non-polar molecules dissolve in non-polar solvents Example: Oils and gasoline Molecular Solids Molecular solids are made of molecule arranged in a pattern that the attractive forces between the molecules. Example (figure on right for ice) Other Important Noncovalent Interactions Noncovalent interactions are interactions that do not involve the sharing of valence electrons such as covalent bonding. Other noncovalent interactions due to the attraction of permanent charges: Salt bridges Ion-dipole interactions A salt bridge is another name for ionic bond that occurs in protein. It is an electrostatic attractive force between a negative formal charge and a positive formal charge in protein. A protein molecule (represented by the ribbon structure to the right) contains positive and negative formal charges. The charges are attracted to each other through salt bridge interactions (indicated by the dashed double arrow). Ion-dipole interactions occur between ions with a full charge and atoms with a partial charge. 28

29 Summary of Noncovalent Interactions Noncovalent Interaction Interaction Between: Hydrogen Bonding Electrostatic attractive force between the partially positive charged hydrogen end of an O-H, N-H, or F-H bond and the negative charge of a lone pair on an O, F, or N. Dipole-Dipole Electrostatic attractive force between two polar molecules. London Dispersion Forces Electrostatic attractive force between any two molecules. Ion-Dipole Electrostatic attractive force between a dipole and an ion or formal charge. The partially positive charge of a dipole is attracted to a negatively charged ion or a negative formal charge. The partially negative charge of a dipole is attracted to a positively charged ion or positive formal charge. Salt Bridge Electrostatic attractive force between a negative formal charge and a positive formal charge in protein. 29

30 An Introduction to Organic Compounds Hydrocarbons Hydrocarbons contain only carbon and hydrogen atoms. There are different types of hydrocarbons: alkanes alkenes alkynes aromatic compounds Alkanes are molecules that consist only of carbon and hydrogen atoms and contain only single bonds. 30

31 Examples of alkanes: Those were examples of alkanes, which means that their carbon chains are unbranched. Properties of Alkanes Alkanes contain only nonpolar covalent bonds, are molecules, and are attracted to one another by London forces. The more carbon atoms in a normal alkane, the higher its point. The longer the alkane, the greater it s surface area and the stronger the London forces that holds it to other molecules. 31

32 Naming Alkanes IUPAC rules, devised by the International Union of Pure and Applied Chemistry, are a widely used method of naming organic compounds. Using the IUPAC rules involves identifying an alkane s (the longest continuous chain of carbon atoms in the molecule) and (atoms or groups of atoms attached to the parent chain). In alkanes, the substituents, called, are constructed solely of carbon and hydrogen atoms. 1. Name the parent chain. The parent chain is named by combining a numbering prefix which specifies the number of carbon atoms in the parent, with, which identifies the molecule as an alkane. 3 carbons = prop ; alkane = ane Propane is an alkane with 3 carbons. 2. Name any alkyl group substituents. Prefixes for Naming Organic Compounds: 32

33 3. Determine the point of attachment of alkyl groups to the parent chain. The parent chain is numbered from the end nearer the alkyl group. 4. Construct the name of the alkane by placing the alkyl groups in alphabetical order and specifying their position numbers, followed by the name of the parent chain 3-ethyl-2-methylheptane The labels di, tri, tetra, etc., are added if two or more identical substituents are present. Note: DO NOT CONSIDER THE PREFIXES (di-, tri-, tetra-) FOR LISTING THE ALKYL GROUPS IN ALPHABETICAL ORDER; ALPHABETIZE BASED ON THE ALKYL GROUP NAME ONLY (methyl, ethyl, propyl, isopropyl, etc). Examples: In the example on the right (5-ethyl-3,3-dimethyloctane), we list ethyl before dimethyl because e comes before m. Give the correct IUPAC name for these molecules. CH 3 CHCH 2 CH 2 CH 2 CH 2 CH 3 CH 3 33

34 Constitutional Isomers Molecules that have the same molecular formula, but different atomic connections are called. Constitutional Isomers of C 6 H 14 Constitutional isomers always have a different name from one another. Group Work: Draw the constitutional isomers for (C 4 H 10 ). Conformations Rotation about allows most molecules to assume a number of different 3-dimensional shapes. The shapes that a molecule can take because of bond rotations are called conformations. The different conformations of a molecule: have the same molecular formula have the same atomic connections have a different 3-dimensional shape 34

35 Conformations of Butane Rotation about the bond between carbons 2 and 3 in butane gives rise to different conformations (three-dimensional shapes) for the molecule. Cycloalkanes In some alkanes, called, carbon atoms are joined into rings. 35

36 Naming Cycloalkanes When naming cycloalkanes, the ring is usually designated as the parent, which is named by combining with the alkane name. Example: When a ring holds more than one alkyl group, the ring is numbered from the position and in the direction that gives the numbers. When there is only one substituent, the 1- is omitted 36

37 When more than one of the same substituent is present on the same carbon, the following method is used: When substituents are non-identical, the ring-carbon attached to substituent that comes first in alphabetical order is assigned position number one. Example: Stereoisomers The rotation of the carbon-carbon single bonds in has an interesting side effect in that it allows for the existence of stereoisomers, molecules that: have the same molecular formula have the same atomic connections have a different 3-dimensional shape are interchanged only by breaking bonds We will study various types of stereoisomers in this course. The first type that we will learn about are called geometric isomers. When stereoisomers exist because of, the stereoisomers are called geometric isomers. 37

38 Geometric Isomers Geometric isomers come in pairs one is cis and one is trans. For cycloalkanes, a cis geometric isomer has two alkyl groups on the same face of the ring and a trans isomer has them on opposite faces. Important NOTE: To visualize geometric isomers (cis and trans), we must draw Side View Structures (as shown above). cis vs. trans cannot be shown using line bond or skeletal structures for cycloalkanes. You try it: Draw and name the geometric isomers of 1-ethyl-2-methylcyclohexane. 38

39 Alkenes, Alkynes, and Aromatic Compounds Saturated vs. Unsaturated Hydrocarbons Saturated hydrocarbons contain only single bonded carbon atoms. alkanes and cycloalkanes Unsaturated hydrocarbons contain double or triple bonds between carbon atoms. alkenes, alkynes, and aromatic compounds Unsaturated Hydrocarbons contain carbon-carbon double bonds. contain carbon-carbon triple bonds. contain benzene rings. Properties of Unsaturated Hydrocarbons Like alkanes, the unsaturated hydrocarbons are nonpolar molecules. London forces hold members of these hydrocarbon families to one another. Increasing molecules size leads to stronger London force attractions and higher melting and boiling points. Naming Unsaturated Hydrocarbons When naming alkenes or alkynes, the parent chain is the longest chain of carbon atoms that the carbon-carbon double or triple bond. 39

40 Numbering the parent begins at the end the double or triple bond. The of the multiple bond is written directly in front of the hydrocarbon parent chain name. Substituents are numbered as in alkanes Naming Aromatic Compounds For structures with benzene, use benzene as the parent name. ortho (o) = 1,2- meta (m) = 1,3- para (p) = 1,4- Polycyclic Aromatic Hydrocarbons (PAHs) 40

41 Geometric Isomerism in Alkenes Free rotations cannot occur in multiple bonds as is does in single bonds. The rotation of the carbon-carbon double bonds in alkenes allows for the existence of cis and trans stereoisomers. Important NOTE: To visualize geometric isomers (cis and trans), we draw all the bonds at 120 o angles from the double bonded carbons! (as shown above). Example: see cis- and trans-2-pentene (above) condensed structure. For geometric isomers to exist, neither carbon atom of the double bond may carry two attached atoms or groups of atoms. Example: Is this the cis or trans geometric isomer of muscalure? 41

42 Functional Groups A is an atom, groups of atoms, or bond that gives a molecule a particular set of chemical properties. Each family of organic compounds is defined by the functional group that contains. All contain a hydroxyl (-OH) functional group that is attached to an alkane-type carbon atom. Example of an alcohol: ethanol contain a carboxyl functional group, which is the combination of a hydroxyl (-OH) group next to a carbonyl group (C=O) group. Examples of carboxylic acids: are carboxylate (CO 2 ) groups between two hydrocarbons. Example of an ester: 42

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