ORGANIC CHEMISTRY 171

1. ORGANIC CHEMISTRY 171 Section 201

2. Chapter one

3. ORIGIN OF ORGANIC CHEMISTRY • In the earlier period of development of chemistry, chemists tried their best to synthesize organic compounds in the laboratory,but failed.Their failures led them to believe that organic compounds could be prepared only from living organisms ( plants and animals ) through force called vital force and theory known as Vital force theory.

4. In 1828 Freidrich Wohler, a German chemist accidentally obtained urea, (NH2)2 CO, an organic compound found in the urine of mammals. In fact, Wohler tried to prepare ammonium cyanate, a substance with mineral origin, by heating ammonium sulphate and potassium cyanate, ammonium cyanate rearranged to urea, a compound which was of organic nature.

5. Freidrich Wohler NH4CNO HEATING NH2CONH2 Ammonium cyanate Urea

6. This chance discovery of Wohler brought about a revolution in the field of organic chemistry. Hennel (1828) prepared ethyl alcohol, Berthelot (1856) prepared methane etc. in the laboratory from mineral resources. By the middle of nineteenth century, the Vital force theory was completely discarded. Chemists then never looked back and at present about ninety five per cent of the organic compounds are man-made.

7. Organic Chemistry • The study of the compounds of carbon. • Over 10 million compounds have been identified. • About 1000 new ones are identified each day! • C is a small atom. • It forms single, double and triple bonds. • It is intermediate in electronegativity (2.5). • It forms strong bonds with C, H, O, N, and some metals.

8. Chemical Bonds Bonds : forces that hold groups of atoms together and make them function as a unit. • ,determines chemical and physical properties of substances

9. Types of Chemical Bonds • Ionic Bond • electrostatic attraction between closely packed oppositely charged ions • formed when a metal reacts with a nonmetal • a metal loses electrons easily to form a positive ion which is attracted to the negative ion formed when the nonmetal gains electrons. • In forming Na+F-, the single 3s electron from Na is transferred to the partially filled valence shell of F.

10. Covalent Bonding • electrons are shared between nuclei • electrons are shared equally between identical atoms. • The simplest covalent bond is that in H2 • polar covalent bond • electrons may be shared unequally between two different atoms. • Or Unequal sharing of electrons between atoms in a molecule.(as HF and HCl) • Results in a charge separation in the bond (partial positive and partial negative charge).

11. Electron Configuration of Atoms • Shells are divided into subshells called orbitals, which are designated by the letters s, p, d, f........ • s (one per shell) • p (set of three per shell 2 and higher) • d (set of five per shell 3 and higher) ..… • The distribution of Orbitals in Shells

12. A molecule is an aggregate of two or more atoms in a definite arrangement held together by chemical bonds H2 H2O NH3 CH4 Molecule, is the collection of two atoms by sharing electrons A diatomic molecule contains only two atoms H2, N2, O2, Br2, HCl, CO A polyatomic molecule contains more than two atoms O3, H2O, NH3, CH4 2.5

13. 11 protons 11 electrons 11 protons 10 electrons Na+ Na 17 protons 18 electrons 17 protons 17 electrons Cl- Cl An ion is an atom, or group of atoms, that has a net positive or negative charge. cation – ion with a positive charge If a neutral atom loses one or more electrons it becomes a cation. anion – ion with a negative charge If a neutral atom gains one or more electrons it becomes an anion. 2.5

14. Electronegativity • Electronegativity • the ability of an atom in a molecule to attract shared electrons to itself. • Pauling scale • Generally increases left to right in a row. • Generally increases bottom to top in a column.

15. Electron Configuration of Atoms • Aufbau (“Build-Up”) Principle: • Orbitals fill in order of increasing energy from lowest energy to highest energy. • Pauli Exclusion Principle: • No more than two electrons may be present in an orbital. If two electrons are present, their spins must be paired. • Hund’s Rule: • one electron is added to each orbital of equal energy before a second electron is added to any one of them; the spins of the electrons should be aligned.

16. Electron Configuration • Energy-level diagram; electrons are placed in an electron configuration. For example, the energy-level diagram for the ground-state electron configuration of carbon is 1s2 2s2 2p2. For chlorine: 1s2 2s2 2p6 3s2 3p5.

17. Lewis Dot Structures • Table 1.4 Lewis Dot Structures for Elements 1-18

18. Formal Charge • Formal charge: The charge on an atom in a molecule or a polyatomic ion. • To derive formal charge 1. Write a correct Lewis structure for the molecule or ion. 2. Assign each atom all its unshared (nonbonding) electrons and one-half its shared (bonding) electrons. 3. Compare this number with the number of valence electrons in the neutral, unbonded atom. 4. The sum of all formal charges is equal to the total charge on the molecule or ion.

19. Formal Charge • Example:Draw Lewis structures, and show which atom in each bears the formal charge.

20. Structural formula of organic compounds The structural formula of a chemical compound is a graphical representation of the molecular structure, showing how the atoms are arranged. The chemical bonding within the molecule is also shown. The structure of carbon and its compound can be expressed using the Lewis-dot structure This system identifies how the atoms in a molecule of a specific compound are attached (bonded) to one another oriented in space.  An atom is indicated by its symbol with a number of dots representing the number of valency electrons e.g Hydrogen would be H with a single dot, Carbon would be a C with four dots

21. When two valency electrons are paired they are represented by two adjacent dots.   Paired valency electrons are not normally available for forming bonds with other atoms.

22. In neutral molecules • hydrogen has one bond. • carbon has 4 bonds and no lone pairs. • nitrogen has 3 bonds and 1 lone pair. • oxygen has 2 bonds and 2 lone pairs. • halogens have 1 bond and 3 lone pairs.

23. Isomers • Two compounds are considered isomers if they have the same molecular formula (i.e. the same numbers and types of atoms) but different structures. • There are two types of isomers, structural isomers and stereoisomers . • structural isomers If two compounds are considered to have the same molecular formula but different connections between atoms (bonding). • stereoisomers If two compounds are considered to have the same molecular formula, the same connections between atoms, but different arrangements of the atoms in three dimensional space. Geometric isomerism (also known as cis-trans isomerism or E-Z isomerism) is a form of stereoisomerism.

24. Consider compounds with the molecular formula C4H10. There are two valid structures that fit this molecular formula.They differ in physical properties. A linear structure: n-butane A branched structure: iso-butane (or 2-methylpropane)

25. Two isomers of butane C4H10: CH3CH2CH2CH3n-butane bp 0 oC mp –138 oC d 0.622 g/cc CH3 CH3CHCH3isobutane bp -12 oC mp -159 oC d 0.604 g/cc

26. Resonance • Linus Pauling - 1930s • Many molecules and ions are best described by writing two or more Lewis structures. • Individual Lewis structures are called contributing structures. • Connect individual contributing structures by double-headed (resonance) arrows. • The molecule or ion is a hybrid of the various contributing structures.

27. Resonance • For many molecules and ions, no single Lewis structure provides a truly accurate representation.

28. Resonance • Examples:equivalent contributing structures.

29. Resonance • Curved arrow:A symbol used to show the redistribution of valence electrons. • In using curved arrows, there are only two allowed types of electron redistribution: • from a bond to an adjacent atom. • from a lone pair on an atom to an adjacent bond.

30. Resonance • All contributing structures must 1. have the same number of valence electrons. 2. obey the rules of covalent bonding: • no more than 2 electrons in the valence shell of H. • no more than 8 electrons in the valence shell of a 2nd period element. 3. differ only in distribution of valence electrons; the position of all nuclei must be the same. 4. have the same number of paired and unpaired electrons.

31. Resonance • The carbonate ion • Is a hybrid of three equivalent contributing structures. • The negative charge is distributed equally among the three oxygens.

32. Resonance • Preference 1: filled valence shells • Structures in which all atoms have filled valence shells contribute more than those with one or more unfilled valence shells.

33. Resonance • Preference 2: maximum number of covalent bonds • Structures with a greater number of covalent bonds contribute more than those with fewer covalent bonds.

34. Resonance • Preference 4: negative charge on the more electronegative atom. • Structures that carry a negative charge on the more electronegative atom contribute more than those with the negative charge on a less electronegative atom.

35. Shapes of Atomic s and p Orbitals • All s orbitals have the shape of a sphere with the center of the sphere at the nucleus.

36. Shapes of Atomic s and p Orbitals • Figure 1.9 (a) 3D representations of the 2px, 2py, and 2pz atomic orbitals including nodal planes.

37. Shapes of Atomic s and p Orbitals • Figure 1.9(b) Cartoon representations of the 2px, 2py, and 2pz atomic orbitals.

38. Hybridization The tetravalency shown by carbon is actually due to excited state of carbon which is responsible for carbon bonding capacity. If the bond formed is by overlapping then all the bonds will not be equivalent so a new concept known as hybridization is introduced which can explain the equivalent character of bonds. s and p orbital belonging to the same atom having slightly different energies mix together to produce same number of new set of orbital called as hybrid orbital and the phenomenon is called as hybridization.

39. Important characteristics of hybridization • The number of hybridized orbital is equal to number of orbitals that get hybridized. • (ii) The hybrid orbitals are always equivalent in energy and shape. • (iii) The hybrid orbitals form more stable bond than the pure atom orbital. • (iv) The hybrid orbitals are directed in space in same preferred direction to have some stable arrangement and giving suitable geometry to the molecule.

40. Types Of Hybridization • There are the three kinds of hybridization that are important in organic chemistry. • (i) sp3 hybridization: In this case, one s and three p orbitalshybridise to form four sp3 hybrid orbitals. These four sp3 hybrid orbitals are oriented in a tetrahedral arrangement.For example in methane CH4

42. (ii) sp2 hybridization: In this case one s and two p orbitals mix together to form three sp2 hybrid orbitals and are orientedin a trigonal planar geometry. Take, for example, ethene (C2H4). Ethene has a double bond between the carbons. For this molecule, carbon will sp2 hybridise, because one π (pi) bond is required for the double bond between the carbons, and only three σ bonds are formed per carbon atom. In sp2 hybridisation the 2s orbital is mixed with only two of the three available 2p orbitals:

43. forming a total of 3 sp2 orbitals with one p-orbital remaining. In ethylene (ethene)or ehtylene the two carbon atoms form a σ bond by overlapping two sp2 orbitals and each carbon atom forms two covalent bonds with hydrogen by s–sp2 overlap all with 120° angles. The π bond between the carbon atoms perpendicular to the molecular plane is formed by 2p–2p overlap. The hydrogen-carbon bonds are all of equal strength and length, which agrees with experimental data.

45. (iii) sp hybridization: In this case, one s and one p orbital mix together to form two sp hybrid orbitals and are oriented in a linear shape. In this model, the 2s orbital mixes with only one of the three p-orbitals resulting in two sp orbitals and two remaining unchanged p orbitals. The chemical bonding in acetylene (ethyne) (C2H2) consists of sp–sp overlap between the two carbon atoms forming a σ bond and two additional π bonds formed by p–p overlap. Each carbon also bonds to hydrogen in a sigma s–sp overlap at 180° angles.