1. Chemical Bonding

2. Chemical Bonds are the forces that hold atoms together in compounds Ionic bondCovalent bondMetallic bond Polar Covalent bond Nonpolar Covalent bond

3. Lewis theory, valence electrons play a fundamental role in chemical bonding e - transfer leads to ionic bonds. Sharing of e - leads to covalent bonds. Atoms tend to have the electron configurations of the noble gas, (octet rule.)

4. valence electrons are the electrons that occupy the outer shell (principal shell number (n)) of an atom. Core electrons are the electrons in inner shells. Lewis symbols : (electron dot symbols) is the simple way of showing the valence electrons of atoms. Si N P As Sb Bi Al Se Ar I

5. Lewis Structures for Ionic Compounds Ba O O Ba 2+ 2- BaO

6. Mg Cl Cl Cl Mg 2+ - 2 MgCl 2

7. Suda çözünürler Yüksek kaynama noktasına sahiptirler - iyonlar arası çekim kuvveti yüksektir Eritilmiş halleri veya sulu çözeltileri elektriği iletir. İyonik bileşikler katı halde elektriği iletmez. İyonik Bileşikler

8. Lewis Structures of covalent bonds Lone pair (nonbonding or unshared) electrons are the electrons are not used for bonding.

10. H N H H H N H H H H + Cl Coordinate Covalent Bonds If both electrons of the bond are contributed by the same atom, this type of bond is called as a coordinate cocalent bond. Cu2+ + 4 NH 3 → Cu(NH 3 ) 4 (type of covalent bond in which both electrons are donated by the same atom

11. En temel yapıtaşları moleküllerdir Erime ve kaynama noktaları düşüktür Elektrik akımını iletmezler Çoğu suda çözünmez Kovalent Bileşikler

12. Multiple bonds In many molecules atoms attain complete octets by sharing more than one pair of electrons between them. The sharing of a pair of electrons represents a single covalent bond, usually just referred to as a single bond, if two electron pairs are shared it is a double bond, if three electron pairs are shared it is called as a triple bond Bond length is the distance between the nuclei of two atoms joined by covalent bond. As a general rule, the distance between bonded atoms decreases as the number of shared electron pairs increases Bond energy (strength) is the quantity of energy required to break one mole of covalent bond. Bond Order –Single bond, order = 1 Double bond, order = 2 Higher bond order –Shorter bond - Stronger bond

15. Bond energy (strength) is the quantity of energy required to break one mole of covalent bond.

16. Bond Energies and Enthalpy of Reaction ΔHrxn= ΔH(product bonds) - ΔH(reactant bonds)

17. = ΔH bonds formed - ΔH bonds broken = -770 kJ/mol – (657 kJ/mol) = -114 kJ/mol

18. Bond polarity In covalent bonds the atoms do not share the electrons equally. The atom which has greater electronegativity, attract the electrons more strongly in covalent bond. nonpolar covalent bondpolar covalent bond electrons are shared equally between two atoms in a chemical bond one atom has a greater attraction for the electrons than the other atom in a chemical bond. Electronegativity: is the relative ability of atoms to attract electrons in a chemical bond.

19. if the electronegativity difference of atoms is 0, the bond is non-polar covalent If the difference in electronegativities between the two atoms is greater than 0, but less than 2.0, the bond is polar covalent, If the difference in electronegativities between the two atoms is 2.0, or greater, the bond is ionic

21. if the electronegativity difference of atoms is 0, the bond is non-polar covalent If the difference in electronegativities between the two atoms is greater than 0, but less than 2.0, the bond is polar covalent, If the difference in electronegativities between the two atoms is 2.0, or greater, the bond is ionic CompoundF2F2 HFLiF Electronegativity Difference 4.0 - 4.0 = 04.0 - 2.1 = 1.94.0 - 1.0 = 3.0 Type of BondNonpolar covalentPolar covalent Ionic (non-covalent)

23. Writing Lewis structure Skeletal structure shows how atoms are attached to one another, the skeletal structure consist of one or more central atoms and terminal atoms. A central atom bonds to two or more atoms in the structure. The atom with the lowest electroneagtivity is generally central atom. Terminal atom bonds to one another atom. Hydrogen atom are terminal atoms.

24. In writing Lewis structure first determine the total number of valence electrons write a possible skeletal structure, connect the atoms by a single covalent bond. (Molecules and polyatomic ions usually have compact, symmetrical structures, but organic compounds are based on long chains of carbon atoms. In oxyacids hydrogen atoms are usually bonded to oxygen atoms) place pairs of electrons as lone pairs around the terminal atoms, according to octet rule. place the remaining electron as lone pairs around the central atom If there are not enough electrons to give the central atom an octet, try multiple bonds (use one or more of the unshared pairs of electrons on the atoms bonded to the central atom to form double or triple bonds, the double bonds form generally among carbon, nitrogen oxygen and sulfur.)

25. question: write a plausible Lewis sturucture for phosgene COCl 2 NO 3 -

26. Formal Charge is the differences between number of valence electrons in a free (uncombined) atom and number of electrons assigned to that atom in the Lewis structure. FC= (number of valence electrons)- ½(number of bonding electrons)-(number of lone pair electrons) -Total formal charges on the atoms in a lewis structure must be 0 to for a neutral atom (and/or to the net charge of a polyatomic ion) - negative formal charges should appear on the most electronegative atoms - where formal charges are required, these should be as small as possible. FC = # valence e- - # lone pair e- - # bond pair e- 2 1

27. question : write a plausible Lewis structure and calculate the formal charge each atom in that formula for each compound given below, NO 3 -, CO 2, NH 4 +, HNO 3, C 2 H 4 SO 4 2- PCl 5, SF 6 NO SOCl 2 ICl 4

28. Resonance: sometimes a molecule or ion can be represented by two or more plausible Lewis structure that differ only by the distribution of electrons. The different plausible structures are called resonance structures. O-O=O  O-O=O

29. Exceptions to the octet rule Odd-electron species NO Incomplete octets BF 3 Expanded octets PCl 5, SF 6

30. H—C—H H O—H B F FF B F FF - + B F FF - + P Cl P Cl Cl Cl Cl S F F F F F F

32. Dipole moments and molecular shape In many cases, covalent bonds are polar covalent because the bound atoms have different electronegativities but whole molecule might be nonploar. The "charge distribution" of a molecule is determined by - The shape of the molecule - The polarity of its bonds A Polar Molecule: The center of the overall negative charge on the molecule does not coincide with the center of overall positive charge on the molecule The molecule can be oriented such that one end has a net negative charge and the other a net positive charge, i.e. the molecule is a dipole A Nonpolar molecule Has no charges on the opposite ends of the molecule Or, has charges of the same sign on the opposite ends of the molecule Molecule is not a dipole question : determine the polarity of following molecules CH4, CH3Cl