1. Organic Chemistry The study of the compounds of carbon Over 10 million compounds have been identified 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

2. Schematic View of an Atom ◦ a small dense nucleus, diameter 10 -14 - 10 -15 m, which contains positively charged protons and most of the mass of the atom ◦ an extranuclear space, diameter 10 -10 m, which contains negatively charged electrons

3. Electron Configuration of Atoms Electrons are confined to regions of space called principle energy levels (shells) ◦ each shell can hold 2n 2 electrons (n = 1,2,3,4......)

4. Electron Configuration of Atoms Aufbau Principle: Aufbau Principle: ◦ orbitals fill in order of increasing energy from lowest energy to highest energy Pauli Exclusion Principle: Pauli Exclusion Principle: ◦ only two electrons can occupy an orbital and their spins must be paired Hund’s Rule: Hund’s Rule: ◦ when orbitals of equal energy are available but there are not enough electrons to fill all of them, one electron is added to each orbital before a second electron is added to any one of them

5. Electron Configuration of Atoms

6. Drawing electron

7. Lewis Dot Structures Gilbert N. Lewis Valence shell: Valence shell: ◦ the outermost occupied electron shell of an atom Valence electrons: Valence electrons: ◦ electrons in the valence shell of an atom; these electrons are used to form chemical bonds and in chemical reactions Lewis dot structure: Lewis dot structure: ◦ the symbol of an element represents the nucleus and all inner shell electrons ◦ dots represent valence electrons

8. Struktur Lewis Dot Struktur Dot Lewis untuk unsur 1-18

9. Lewis Model of Bonding Atoms bond together so that each atom acquires an electron configuration the same as that of the noble gas nearest it in atomic number anion ◦ an atom that gains electrons becomes an anion cation ◦ an atom that loses electrons becomes a cation ionic solids ◦ the attraction of anions and cations leads to the formation of ionic solids covalent bond ◦ an atom may share electrons with one or more atoms to complete its valence shell; a chemical bond formed by sharing electrons is called a covalent bond polar covalent bonds ◦ bonds may be partially ionic or partially covalent; these bonds are called polar covalent bonds

10. In chemistry, a bond is typically classified as one of two types: covalent Purely covalent (non-polar): The bonding electrons are shared equally between the two bonding atoms. Polar covalent: The electrons are shared between the two bonding atoms, but unequally, with the electrons spending more time around the more electronegative atom. Ionic Ionic: The electrons aren’t shared. Instead, the more electronegative atom of the two bonding atoms selfishly grabs the two electrons for itself, giving this more electronegative atom a formally negative charge and leaving the other atom with a formal positive charge. The bond in an ionic bond is an attraction of opposite charges.

11. Electronegativity Electronegativity: Electronegativity: ◦ a measure of an atom’s attraction for the electrons it shares with another atom in a chemical bond Pauling scale Pauling scale ◦ generally increases left to right in a row ◦ generally increases bottom to top in a column

12. Electronegativity: The ability of atoms in a molecule to attract electrons to itself. On the periodic chart, electronegativity increases as you go… ◦ …from left to right across a row. ◦ …from the bottom to the top of a column. Electronegativity increases Electronegativity increases

13. Formation of Ions A rough guideline: ◦ ions will form if the difference in electronegativity between interacting atoms is 1.9 or greater ◦ example: sodium (EN 0.9) and fluorine (EN 4.0) ◦ we use a single-headed (barbed) curved arrow to show the transfer of one electron from Na to F ◦ in forming Na + F -, the single 3s electron from Na is transferred to the partially filled valence shell of F

14. Covalent Bonds The simplest covalent bond is that in H 2 ◦ the single electrons from each atom combine to form an electron pair ◦ the shared pair functions in two ways simultaneously; it is shared by the two atoms and fills the valence shell of each atom The number of shared pairs ◦ one shared pair forms a single bond ◦ two shared pairs form a double bond ◦ three shared pairs form a triple bond

15. Polar and Non-polar Covalent Bonds Although all covalent bonds involve sharing of electrons, they differ widely in the degree of sharing We divide covalent bonds into ◦ Non-polar covalent bonds ◦ polar covalent bonds

16. Polar and Non-polar Covalent Bonds ◦ an example of a polar covalent bond is that of H-Cl ◦ the difference in electronegativity between Cl and H is 3.0 - 2.1 = 0.9  +  - ◦ we show polarity by using the symbols  + and  -, or by using an arrow with the arrowhead pointing toward the negative end and a plus sign on the tail of the arrow at the positive end

17. Polar Covalent Bonds Bond dipole moment (  ): Bond dipole moment (  ): ◦ a measure of the polarity of a covalent bond ◦ the product of the charge on either atom of a polar bond times the distance between the nuclei ◦ Table below shows average bond dipole moments of selected covalent bonds

18. Lewis Structures To write a Lewis structure ◦ determine the number of valence electrons ◦ determine the arrangement of atoms ◦ connect the atoms by single bonds ◦ arrange the remaining electrons so that each atom has a complete valence shell ◦ show a bonding pair of electrons as a single line ◦ show a nonbonding pair of electrons as a pair of dots ◦ in a single bond atoms share one pair of electrons, in a double bond they share two pairs of electrons, and in a triple bond they share three pairs of electrons

19. Lewis Structures 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

20. Formal Charge 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

21. Structure and Bonding

22. Formal Charge Draw Lewis structures, and show which atom in each bears the formal charge ClO 2 - CH 3 COO - HCOO - BH 4 - CH 3 NH 2 C3H4C3H4

23. Exceptions to the Octet Rule Molecules containing atoms of Group 3A elements, particularly boron and aluminum Aluminum chloride : : : FB F F ClAl Cl Cl 6 electrons in the valence shells of boron and aluminum Boron trifluoride :: :: :: : : : : : : : : :

24. Exceptions to the Octet Rule Atoms of third-period elements have 3d orbitals and may expand their valence shells to contain more than 8 electrons ◦ phosphorus may have up to 10

25. Exceptions to the Octet Rule ◦ sulfur, another third-period element, forms compounds in which its valence shell contains 8, 10, or 12 electrons

26. VSEPR Based on the twin concepts that ◦ atoms are surrounded by regions of electron density ◦ regions of electron density repel each other

27. VSEPR Model Example: Example: predict all bond angles for these molecules and ions

28. Polar and Nonpolar Molecules To determine if a molecule is polar, we need to determine ◦ if the molecule has polar bonds ◦ the arrangement of these bonds in space Molecular dipole moment (  ): Molecular dipole moment (  ): the vector sum of the individual bond dipole moments in a molecule ◦ reported in debyes (D)

29. Polar and Nonpolar Molecules these molecules have polar bonds, but each has a zero dipole moment

30. Polar and Nonpolar Molecules these molecules have polar bonds and are polar molecules

31. Polar and Nonpolar Molecules ◦ formaldehyde has polar bonds and is a polar molecule