1. Today’s Quiz 1 1.What is ground-state electron configuration? 2.Define valence electrons and valence shell. 3.Explain the exceptions to the octet rule. 4.Define Electronegativity. 5.What is the difference between a nonpolar covalent bond and a polar covalent bond. 6.What are the 3 molecular shapes? 7.What is a hybrid orbital?

2. 2 Chapter 1: Covalent Bonding and Shapes of Molecules

3. 3 1.1 How do we describe the electronic structure of Atoms? 1.2 What is the Lewis Model of Bonding? 1.3 How do we Predict Bond Angles and the Shapes of Molecules? 1.4 How do we predict if a molecule is polar or non-polar? 1.5 What is resonance? 1.6 What is the Molecular orbital model of covalent bonding? 1.7 What are the Functional Groups? (Group Presentation) Chapter 1: Covalent Bonding and Shapes of Molecules

4. Quantum Numbers and Atomic Orbitals An atomic orbital is specified by four quantum numbers. n the principal quantum number - a positive integer, indicates the relative size of the orbital or the distance from the nucleus l the angular momentum quantum number - an integer from 0 to n-1, related to the shape of the orbital m l the magnetic moment quantum number - an integer from - l to + l, orientation of the orbital in the space around the nucleus 2 l + 1 Schrodinger Wave Equation m s spin quantum number

5. Table 7.2 The Hierarchy of Quantum Numbers for Atomic Orbitals Name, Symbol (Property) Allowed ValuesQuantum Numbers Principal, n (size, energy) Angular momentum, l (shape) Magnetic, m l (orientation) Positive integer (1, 2, 3,...) 0 to n-1 - l,…,0,…,+ l 1 0 0 2 01 0 3 0 12 0 0 +1 0+1 0 +2-2

6.  = fn(n, l, m l, m s ) spin quantum number m s m s = +½ or -½ Schrodinger Wave Equation m s = -½m s = +½ 7.6

7. Electron configuration is how the electrons are distributed among the various atomic orbitals in an atom. 1s 1 principal quantum number n angular momentum quantum number l number of electrons in the orbital or subshell Orbital diagram H 1s 1 7.8

8. Order of orbitals (filling) in multi-electron atom 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s 7.7

9. Ionic Bond Is the electrostatic force that holds ions together in an ionic compound The metal gives the electrons to the non metal Cation is the metal Anion is the non metal

11. The Covalent Bond A bond in which two electrons are shared by two atoms. Only in covalent compounds Non metal and non metal

12. The Covalent Bond

13. Pairs of valence electrons that are not involved in covalent bond formation are called: Lone pairs We can only draw Lewis structures for compounds that have covalent bonds

14. Lewis Dot Symbols Consists of the symbol of an element and one dot for each valence electron in an atom of the element. Note that (except helium) the number of valance electrons each atom has is the same as the group number of the element.

15. Lewis Structures Is a representation of covalent bonding in which shared electron pairs are shown either as lines or as pairs of dots between two atoms, and lone pairs of dots on individual atoms Only valence electrons are shown in a Lewis structure Valence electrons: Electrons in the valence (outermost) shell of an atom. Valence shell: The outermost electron shell of an atom.

16. Electronegativity The ability of an atom to attract toward itself the electrons in a chemical bond

17. Electronegativity

18. 18 I. Lewis Structures B. Ionic, covalent, and polar bonds H—H  = 2.1 2.1  = 0  equal sharing of electrons Cl—Cl= nonpolar covalent bond  = 3.0 3.0 H—Cl  = 0.9  unequal sharing of electrons  = 2.1 3.0 = polar covalent bond Na + Cl –  = 2.1  transfer of electrons  = 0.9 3.0 = ionic bond generally: when  <0.5 non-polar covalent  = 1.9  polar covalent > 1.9  ionic  +  – nonmetal + nonmetal metal + nonmetal

19. Polar and Nonpolar Molecules A molecule will be polar if: 1. It has polar bonds. 2. The center of partial positive charge lies at a different place within the molecule than the center of partial negative charge.

20. The "best" Lewis structure for NO 3 - 1. Determine the total number of valence electrons in a molecule Draw a skeleton

21. Of the 24 valence electrons in NO 3 -, 6 were required to make the skeleton. Consider the remaining 18 electrons and place them so as to fill the octets of as many atoms as possible

22. Are the octets of all the atoms filled? If not then fill the remaining octets by making multiple bonds Check that you have the lowest FORMAL CHARGES possible for all the atoms, without violating the octet rule; (valence e - ) - (1/2 bonding e - ) - (lone electrons).

23. The Octet Rule

24. Exceptions to the Octet Rule The expanded octet – Atoms in the 2 nd period (that are in families 3A-7A) cannot have more than 8 valance electrons around the central atom – Atoms of elements in and beyond the 3 rd period (that are in families 3A-7A) form some compound in which more than 8 electron surround the central atom Odd- electron molecules – Some molecules contain an odd number of electrons. – We need an even number of electron for complete pairing the octet rule clearly cannot be satisfied with all the atoms in any of these molecules Incomplete octet – The number of electrons surrounding the central atom in a stable molecule is fewer than eight

25. Figure 2.10 The modern periodic table. +1 +2 +3-3-2 0 NC

26. The Concept of Resonance Resonance means the use of two or more Lewis structures to represent a particular molecule Resonance structure, is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure

27. 27 E. Resonance structures -two or more equivalent Lewis structures -nuclei remain in fixed positions, but electrons arranged differently neither of these accurately describes the formate ion actual species is an average of the two (resonance hybrid) Resonance hybrid: A molecule or ion that is best described as a composite of a number of contributing structures. delocalized electrons

28. 28 I. Lewis Structures E. Resonance structures more stable major contributor less stable minor contributor

29. Shapes of 2p x, 2p y, 2p z atomic orbitals and their orientation in space

30. tetrahedral Trigonal planar or pyramidal linear Region of electron density around the central atom.

31. Sigma bond is covalent bond in which the overlap of atomic orbitals is concentrated along the bond axis Hybrid orbital an orbital produced from the combination of two or more atomic orbitals. 31 Pi bond a covalent bond formed by the overlap of parallel p orbitals.

32. Why does hybridization occur? In chemistry hybridization is the concept of mixing atomic orbitals to form new hybrid orbitals suitable for the qualitative description of atomic bonding properties. Hybridized orbitals are very useful in the explanation of the shape of molecular orbitals for molecules. It is an integral part of organic chemistry. 32

33. Sp 3 hybrid orbitals – Bond angles of approximately 109.5 o Four sigma bonds 4 groups bonded to carbon Sp 2 hybrid orbitals – Bond angles of approximately 120 o Three sigma bonds and one pi bond 3 groups bonded to carbon Sp hybrid orbitals – Bond angles of approximately 180 o Two sigma bonds and two pi bonds 2 groups bonded to carbon tetrahedral Trigonal planar linear

34. 34 III. Valence Bond Model A. Hybrid atomic orbitals 1. sp 3 hybridization CH 4 facts: tetrahedral, 4 equivalent bonds

35. 35 III. Valence Bond Model B. Hybrid atomic orbitals 1. sp 3 hybridization

36. 36 III. Valence Bond Model B. Hybrid atomic orbitals 2. sp 2 hybridization C 2 H 4 facts: all six atoms lie in same plane trigonal planar = sp 2

37. 37 III. Valence Bond Model B. Hybrid atomic orbitals 2. sp 2 hybridization

38. 38 III. Valence Bond Model B. Hybrid atomic orbitals 3. sp hybridization C 2 H 2 facts:linear = sp

39. 39 III. Valence Bond Model B. Hybrid atomic orbitals 3. sp hybridization