Lec 1,2_Organic 1_Peter PDF

Organic Chemistry (1) First Level Clinical Pharmacy Program Introduction to Organic Chemistry Peter S. Ayoub, PhD. What Is Organic Chemistry? Organic chemistry is the chemistry of compounds that contain the element carbon. Pro...

Organic Chemistry (1) First Level Clinical Pharmacy Program Introduction to Organic Chemistry Peter S. Ayoub, PhD. What Is Organic Chemistry? Organic chemistry is the chemistry of compounds that contain the element carbon. Products of organic chemistry used in medicine What are the common features of these organic compounds? All carbon atoms have four bonds. A All organic compounds contain carbon stable carbon atom is said to be atoms and most contain hydrogen atoms. tetravalent. Other elements may also be present. Any Some compounds have chains of atom that is not carbon or hydrogen is atoms and some compounds have called a heteroatom. Common heteroatoms rings. include N, O, S, P, and the halogens.  The Periodic Table All matter is composed of the same building blocks called atoms. There are two main components of an atom. The nucleus contains positively charged protons and uncharged neutrons. Most of the mass of the atom is contained in the nucleus. The electron cloud is composed of negatively charged electrons. The electron cloud comprises most of the volume of the atom  The Periodic Table All matter is composed of the same building blocks called atoms. There are two main components of an atom. The charge on a proton is equal in magnitude but opposite in sign to the charge on an electron.  The Periodic Table A cation is positively charged and has fewer electrons than its neutral form. An anion is negatively charged and has more electrons than its neutral form.  The Periodic Table Although more than 100 elements exist, most are not common in organic compounds. Most of these elements are located in the first and second rows of the periodic table.  The Periodic Table Across each row of the periodic table, electrons are added to a particular shell of orbitals around the nucleus. The shells are numbered 1, 2, 3, and so on. Adding electrons to the first shell forms the first row. Adding electrons to the second shell forms the second row. Electrons are first added to the shells closest to the nucleus. These electrons are held most tightly.  The Periodic Table Each shell contains a certain number of subshells called orbitals. An orbital is a region of space that is high in electron density. There are four different kinds of orbitals, called s, p, d, and f. The first shell has only one orbital, called an s orbital. The second shell has two kinds of orbitals, s and p, and so on. Each type of orbital occupies a certain space and has a particular shape.  The Periodic Table For the first and second row elements, we must deal with only s orbitals and p orbitals. An s orbital has a sphere of electron density. It is lower in energy than other orbitals of the same shell, because electrons are kept close to the positively charged nucleus. An s orbital is filled with electrons before a p orbital in the same shell.  The Periodic Table For the first and second row elements, we must deal with only s orbitals and p orbitals. A p orbital has a dumbbell shape. It contains a node of electron density at the nucleus. A node means there is no electron density in this region. A p orbital is higher in energy than an s orbital (in the same shell) because its electron density is farther away from the nucleus. A p orbital is filled with electrons only after an s orbital of the same shell is full.  The Periodic Table The outermost electrons are called valence electrons. The valence electrons are more loosely held than the electrons closer to the nucleus, and as such, they participate in chemical reactions. The group number of a second-row element reveals its number of valence electrons. For example, carbon in group 4A has four valence electrons, and oxygen in group 6A has six.  Bonding Bonding is the joining of two atoms in a stable arrangement. Bonding may occur between atoms of the same or different elements. Bonding is a favorable process because it always leads to lowered energy and increased stability. Joining two or more elements forms compounds. Although only about 100 elements exist, more than 30 million compounds are known.  Bonding Examples of compounds include hydrogen gas (H2), formed by joining two hydrogen atoms, and methane (CH4), the simplest organic compound, formed by joining a carbon atom with four hydrogen atoms.  Bonding Examples of compounds include hydrogen gas (H2), formed by joining two hydrogen atoms, and methane (CH4), the simplest organic compound, formed by joining a carbon atom with four hydrogen atoms.  Lewis Structures Lewis structures are electron dot representations for molecules. There are three general rules for drawing Lewis structures. Draw only the valence electrons. Give every second-row element no more than eight electrons. Give each hydrogen two electrons. Lewis Structures How to draw Lewis Structure How to draw Lewis Structure How to draw Lewis Structure How to draw Lewis Structure How to draw Lewis Structure  How to draw Lewis Structure Formal Charge Formal charge is the charge assigned to individual atoms in a Lewis structure. How to draw Lewis Structure  Exceptions to the Octet Rule Elements in groups 2A and 3A of the periodic table, such as beryllium and boron, do not have enough valence electrons to form an octet in a neutral molecule. Lewis structures for BeH2 and BF3 show that these atoms have only four and six electrons, respectively, around the central atom. There is nothing we can do about this! There simply aren’t enough electrons to form an octet. Because the Be and B atoms each have less than an octet of electrons, these molecules are highly reactive.  Exceptions to the Octet Rule A second exception to the octet rule occurs with some elements located in the third row and later in the periodic table. These elements have empty d orbitals available to accept electrons, and thus they may have more than eight electrons around them. For organic chemists, the two most common elements in this category are phosphorus and sulfur, which can have 10 or even 12 electrons around them.  Resonance Resonance structures are two Lewis structures having the same placement of atoms but a different arrangement of electrons.  Resonance ? ? Which resonance structure is an accurate representation for (HCONH)–? The answer is neither of them. The true structure is a composite of both ? ? resonance forms, and is called a resonance hybrid. The hybrid shows characteristics of both Question? resonance structures.  Resonance ? ? Which resonance structure is an accurate representation for (HCONH)–? Each resonance structure implies that electron pairs are localized in bonds or on atoms. ? ? In actuality, resonance allows certain electron pairs to be delocalized over two or more atoms, and this delocalization of electron Question? density adds stability. A molecule with two or more resonance structures is said to be resonance stabilized.  Resonance Basic principles of resonance theory Resonance structures are not real. An individual resonance structure does not accurately represent the structure of a molecule or ion. Only the hybrid does. Resonance structures are not in equilibrium with each other. There is no movement of electrons from one form to another. Resonance structures are not isomers. Two isomers differ in the arrangement of both atoms and electrons, whereas resonance structures differ only in the arrangement of electrons.  Resonance Basic principles of resonance theory  Resonance To draw resonance structures, use the three rules that follow:  Resonance To draw resonance structures, use the three rules that follow:  Resonance To draw resonance structures, use the three rules that follow:  Resonance To draw resonance structures, use the three rules that follow: Resonance structures A and B differ in the location of two electron pairs, so two curved arrows are needed. To convert A to B, take the lone pair on N and form a double bond between C and N. Then, move an electron pair in the C – O double bond to form a lone pair on O. Curved arrows thus show how to reposition the electrons in converting one resonance form to another. The electrons themselves do not actually move. Resonance Resonance Resonance  The Resonance Hybrid The resonance hybrid is the composite of all possible resonance structures. In the resonance hybrid, the electron pairs drawn in different locations in individual resonance structures are delocalized. The resonance hybrid is more stable than any resonance structure because it delocalizes electron density over a larger volume.  The Resonance Hybrid When two resonance structures are different, the hybrid looks more like the “better” resonance structure. The “better” resonance structure is called the major contributor to the hybrid, and all others are minor contributors. The hybrid is the weighted average of the contributing resonance structures. What makes one resonance structure “better” than another? The Resonance Hybrid The Resonance Hybrid  Determining Molecular Shape We can now use Lewis structures to determine the shape around a particular atom in a molecule. Consider the H2O molecule. The Lewis structure tells us only which atoms are connected to each other, but it implies nothing about the geometry. What does the overall molecule look like? Is H2O a bent or linear molecule? Two variables define a molecule’s structure: bond length and bond angle.  Determining Molecular Shape Two variables define a molecule’s structure: Bond Length Bond Angle  Determining Molecular Shape Two variables define a molecule’s structure: Bond Length  Determining Molecular Shape Two variables define a molecule’s structure: Bond Length  Determining Molecular Shape Two variables define a molecule’s structure: Bond Angle Bond angle determines the shape around any atom bonded to two other atoms. To determine the bond angle and shape around a given atom, we must first determine how many groups surround the atom. A group is either an atom or a lone pair of electrons. Then we use the valence shell electron pair repulsion (VSEPR) theory to determine the shape.  VSEPR is based on the fact that electron pairs repel each other; thus: The most stable arrangement keeps these groups as far away from each other as possible.  Determining Molecular Shape Two variables define a molecule’s structure: Bond Angle  Determining Molecular Shape Two variables define a molecule’s structure: Bond Angle  Determining Molecular Shape Two variables define a molecule’s structure: Bond Angle  Determining Molecular Shape Two variables define a molecule’s structure: Bond Angle  Determining Molecular Shape Two variables define a molecule’s structure: Bond Angle  Determining Molecular Shape Two variables define a molecule’s structure: Bond Angle  Determining Molecular Shape Two variables define a molecule’s structure: Bond Angle  Determining Molecular Shape Two variables define a molecule’s structure: Bond Angle  Determining Molecular Shape Two variables define a molecule’s structure: Bond Angle  Determining Molecular Shape Two variables define a molecule’s structure: Bond Angle  Drawing Organic Structures Drawing organic molecules presents a special challenge. Because they often contain many atoms, we need shorthand methods to simplify their structures. The two main types of shorthand representations used for organic compounds are: Condensed structures Skeletal structures  Drawing Organic Structures Condensed structures  Drawing Organic Structures Condensed structures  Drawing Organic Structures Condensed structures containing a C=O bond  Drawing Organic Structures Skeletal structures  Drawing Organic Structures Skeletal structures  Drawing Organic Structures Skeletal structures Drawing Organic Structures  Drawing Organic Structures When heteroatoms are bonded to a carbon skeleton, the heteroatom is joined directly to the carbon to which it is bonded, with no H atoms in between. Thus, an OH group is drawn as OH or HO depending on where the OH is located. In contrast, when carbon appendages are bonded to a carbon skeleton, the H atoms will be drawn to the right of the carbon to which they are bonded regardless of the location.  Drawing Organic Structures Skeletal structures  Drawing Organic Structures Skeletal structures Hybridization Hybridization Hybridization Hybridization Hybridization Hybridization Hybridization Hybridization Hybridization Hybridization Hybridization Hybridization Hybridization  Electronegativity and Bond Polarity Electronegativity is a measure of an atom’s attraction for electrons in a bond. Thus, electronegativity indicates how much a particular atom “wants” electrons.  Electronegativity and Bond Polarity Usually, a polar bond will be one in which the electronegativity difference between two atoms is ≥ 0.5 units. Problem Vs Solution Problem Problem Vs Solution Solution  Assignment ? ? Provide the following about L- Doba. ? Two polar and two nonpolar bonds ? Label all sp3 Hybridized C atoms ? ? ? Label all H atoms that bear a partial positive charge ? Draw another resonance structure Question?  Functional Groups A functional group is an atom or a group of atoms with characteristic chemical and physical properties. It is the reactive part of the molecule.  Functional Groups A functional group is an atom or a group of atoms with characteristic chemical and physical properties. It is the reactive part of the molecule. Functional Groups Functional Groups Functional Groups  Intermolecular Forces Intermolecular forces are the interactions that exist between molecules. A functional group determines the type and strength of these interactions. Covalent compounds are composed of discrete molecules. The nature of the forces between the molecules depends on the functional group present.  Intermolecular Forces There are three different types of inter- actions, presented here in order of increasing strength: Van der Waals forces Dipole–dipole interactions Hydrogen bonding  Intermolecular Forces Van der Waals forces van der Waals forces, also called London forces, are very weak interactions caused by the momentary changes in electron density in a molecule. van der Waals forces are the only attractive forces present in nonpolar compounds. For example, although a nonpolar CH4 molecule has no net dipole, at any one instant its electron density may not be completely symmetrical, creating a temporary dipole. This can induce a temporary dipole in another CH4 molecule, with the partial positive and negative charges arranged close to each other. The weak interaction of these temporary dipoles constitutes van der Waals forces. All compounds exhibit van der Waals forces.  Intermolecular Forces Van der Waals forces  Intermolecular Forces Van der Waals forces  Intermolecular Forces Van der Waals forces  Intermolecular Forces Van der Waals forces  Intermolecular Forces Van der Waals forces  Intermolecular Forces Dipole–Dipole Interactions Dipole–dipole interactions are the attractive forces between the permanent dipoles of two polar molecules. In acetone, (CH3)2C – O, for example, the dipoles in adjacent molecules align so that the partial positive and partial negative charges are in close proximity. These attractive forces caused by permanent dipoles are much stronger than weak van der Waals forces.  Intermolecular Forces Hydrogen Bonding Hydrogen bonding typically occurs when a hydrogen atom bonded to O, N, or F, is electro- statically attracted to a lone pair of electrons on an O, N, or F atom in another molecule. Thus, H2O molecules can hydrogen bond to each other. When they do, an H atom covalently bonded to O in one water molecule is attracted to a lone pair of electrons on the O in another water molecule. Hydrogen bonds are the strongest of the three types of intermolecular forces, though they are still much weaker than any covalent bond.  Intermolecular Forces Hydrogen Bonding Intermolecular Forces Intermolecular Forces  A Picture Is Worth a Thousand Words Review the Concepts Thanks!