Lecture_1_Introduction_to_Organic_Chemistry_&_Structure_of_Organic.pptx

Lecture number: 1 Introduction to Organic Chemistry & Structure of Organic Molecules College of Pharmacy you can find this lect...

Lecture number: 1 Introduction to Organic Chemistry & Structure of Organic Molecules College of Pharmacy you can find this lecture in Essentials of Organic Chemistry book by Paul M Dewick from page 19-53 Department of Pharmaceutical Section 1. Atomic Structure 2. Bonding and Valency 3. Atomic orbitals 4. Electronic configuration 5. Hybrid orbitals in carbon 7. Bond polarity 8. Conjugation 8. Resonance Structures and Curvy Arrows 9. Hydrogen Bonding References 1. Atomic Structure (page 19) Organic chemistry is the study of carbon compounds. Atoms are composed of protons, neutrons and electrons. Protons are positively charged, electrons carry a negative charge, and neutrons are uncharged. In a neutral atom, the nucleus of protons and neutrons is surrounded by electrons, the number of which is equal to the number of protons (atomic number). 9/8/24 3 Continued… If the number of electrons and protons is not equal, the atom or molecule containing the atom will necessarily carry a charge, and is called an ion. A negatively charged atom or molecule is termed an anion, and a positively charged species is called a cation. Acquiring a noble gas-like complement of electrons governs the bonding together of atoms to produce molecules, it can be achieved by losing electrons, by gaining electrons, or by sharing electrons associated with the unfilled shell, and leads to ionic bonds or 9/8/24 4 2. Bonding and Valency (page 19) Ionic bonding. This involves loss of an electron from one atom, and its transfer to another. Covalent Bond: Sharing of electrons effectively brings each atom up to the noble gas electronic configuration. Hydrogen Chloride Hydrogen 9/8/24 chloride 5 Continued… The unfilled shell involved in bonding is termed as the valence shell, and the electrons in it are termed as valence electrons. The number of electrons in the valence shell is equal to the group number of the atom. For example, carbon is in group IVA and carbon has four valence electrons; oxygen is in group VIA and has six valence electrons. The halogens of group VIIA all have seven electrons. 9/8/24 6 3. Atomic orbitals (page 20) Atomic orbitals describe the probability of finding a given electron of an atom in a given region of space. We are unable to pin-point the electron at any particular time, but we have an indication that it will be within certain spatial limits. 9/8/24 7 Continued… Figure: Shapes of atomic orbitals 9/8/24 8 Continued… Figure: Relative energies of atomic orbitals (not to scale) 9/8/24 9 Continued… Figure. Electronic configurations: energy diagrams 9/8/24 10 4. Electronic configuration (page 23) The electronic configuration can be expressed as a list of those orbitals containing electrons. Each atomic orbital can accommodate just two electrons. Electrons are allocated to atomic orbitals, one at a time, so that orbitals of one energy levels are filled before proceeding to the next higher level. When electrons are placed in orbitals of the same energy (degenerate orbitals, e.g. p orbitals) they are located singly in separate orbitals before two electrons are paired. 9/8/24 11 Continued… 9/8/24 12 5. Hybrid orbitals in carbon (page 26) The hydrocarbon methane (CH4) is tetrahedral in shape with bond angles of about 109◦, and the four C–H bonds are all equivalent and identical in reactivity. Ethylene (ethene, C2H4) is planar, with bond angles of about 120◦, and it contains one π bond. Acetylene (ethyne, C2H2) is linear, i.e. bond angles 180◦, and it contains two π bonds. None of these observations follows immediately from the electronic configuration of carbon (1s22s22px12py1), which shows that carbon has two unpaired electrons, each in a 2P orbital. 9/8/24 13 sp3 hybrid orbitals The ability of carbon to bond to four other atoms requires unpairing of the 2s2 electrons. We might consider promoting one electron from a 2s orbital to the third, as yet unoccupied, 2p orbital. We now have four unpaired electrons in separate orbitals, and the electronic configuration of carbon has become1s22s12px12py12p1z. 9/8/24 14 Continued… The hybrid orbitals for carbon are derived by mixing the one 2s orbital and three 2p atomic orbitals. This generates four equivalent hybrid orbitals, which we designate sp3, since they are derived from one s orbital and three p orbitals. The sp3 orbitals will be at an energy level intermediate between those of the 2s and 2p orbitals. 9/8/24 15 sp2 hybrid orbitals To provide a model for ethylene, we now need to consider hybrid orbitals that are a mix of the 2s orbital with two 2p orbitals, giving three equivalent sp2 orbitals. The energy level associated with an sp2 orbital will be below that of the sp3 orbital: this time, we have mixed just two high-energy p orbitals with the lower energy s orbital. The 2 p orbital is then located perpendicular to this plane. 9/8/24 16 Continued… Figure. Electronic configuration: sp2-hybridized carbon atom 9/8/24 17 sp hybrid orbitals The third observation relates to acetylene (ethyne, C2H2), which is linear, i.e. bond angles of 180◦, and contains two π bonds. This introduces what we term triple bonds, actually a combination of one σ bond and two π bonds. In this molecule, we invoke another type of hybridization for carbon, that of sp hybrid orbitals. The bonding in acetylene has one C–C σ bond together with two C–H σ bonds; the p orbitals on each carbon, each carrying one electron, interact by side-to-side overlap to produce two π bonds. Note that the p orbitals can only overlap if their axes are parallel. 9/8/24 18 Continued… Figure. Electronic configuration: sp-hybridized carbon atom 9/8/24 19 6. Bond polarity (page 35) The nucleus of each atom has a certain ability to attract electrons. This is termed its electronegativity. This means that, when it is bonded to another atom, the bonding electrons are not shared equally between the two atoms. This is indicated in a structure by putting partial charges (δ+ and δ−) above the atoms. 9/8/24 20 7. Conjugation (page 37) We use the term conjugated to describe an arrangement in which double bonds are separated by a single bond. 9/8/24 21 8. Resonance Structures and Curvy Arrows (page 45) The curly arrow represents the movement of two electrons. The tail of the arrow indicates where the electrons are coming from, and the arrowhead where they are going to. Curly arrows must start from an electron-rich species. This can be a negative charge, a lone pair, or a bond. Arrowheads must be directed towards an electron deficient species. This can be a positive charge, the positive end of a polarized bond, or a suitable atom capable of accepting electrons, i.e. an electronegative atom. Examples for resonance cases: 1- conjugated systems 2- Double beside oxygen 3- Double bond beside nitrogen 9/8/24 22 Continued… The formal charge on an individual atom can be assessed more rigorously by subtracting the number of valence electrons assigned to an atom in its bonded state from the number of valence electrons it has as a neutral free atom. Electrons in bonds are considered as shared equally between the atoms, whereas unshared lone pairs are assigned to the atom that possesses them. Formal charge = number of valence electrons number of valence electrons as neutral free atom – assigned in bonded state 9/8/24 23 Continued… 9/8/24 24 9. Hydrogen bonding (page 49) Hydrogen bonds (H-bonds) describe the weak attraction of a hydrogen atom bonded to an electronegative atom, such as oxygen or nitrogen, to the lone pair electrons of another electronegative atom. These bonds are different in nature from the covalent bonds. 9/8/24 25 References Essential of Organic Chemistry for students of pharmacy, medicinal chemistry, and biological chemistry, Paul M Dewick, Edition 2006. Pages 19-53. 9/8/24 26