1. Covalent Bonding

2. Lesson 1:Covalent Bonding Covalent bonds: atoms held together by sharing electrons. Mostly formed between nonmetals Molecules: neutral group of atoms joined together by covalent bonds. Diatomic molecule: molecule consisting of 2 atoms. Remember them: F 2, Cl 2, I 2, Br 2, H 2, N 2, O 2 Molecules tend to have lower melting and boiling points than ionic compounds.

3. The Nature of Covalent Bonding Introduction with balloon activity octet rule: electron sharing occurs usually so that atoms attain the electron configurations of noble gases. Single covalent bond: two atoms held together by sharing a pair of electrons. Shown as two dots or as a long dash. A pair of valence electrons that is not shared between atoms is called an unshared pair.

4. HH O H O H or O H H

5. Double bonds: covalent bond formed by sharing two pairs of electrons Triple bonds: covalent bond formed by sharing three pairs of electrons.

6. hydrogen chlorine iodine nitrogen Covalent bonding with equal sharing of electrons occurs in diatomic molecules formed from one element. A dash may replace a pair of dots.

7. Strength of covalent bonds Covalent bonds differ in strength, some break more easily than others. Bond length: the shorter the bond length, the stronger the bond. Triple bond is stronger than a double bond; double bond is stronger than a single bond. Bonds and energy: Energy is released when a bond forms; energy is added to break a bond. The smaller the bond length, the greater the bond- dissociation energy (harder to break bond).

8. Classwork Read section 8.1 in textbook and do Section 8.1 assessment

9. Lesson 2: Molecular Formula Shows how many atoms of each element a molecule contains. Naming binary molecular compounds Some molecules are known for their common names: Ex. Ammonia NH 3 Composed of two nonmetals; often combine in more than one way. Ex. CO and CO 2 Greek Prefixes are used to name binary molecular compounds. PrefixMono-Di-Tri-Tetra-Penta-Hexa-Hepta-Octa-Nona-Deca- Number 12345678910

10. Binary Compounds Containing Two Nonmetals To name these compounds: 1)give the name of the less electronegative element first with the Greek prefix indicating the number of atoms of that element present 2) After give the name of the more electronegative non- metal with the Greek prefix indicating the number of atoms of that element present and with its ending replaced by the suffix –ide. 3)Do not use the prefix mono- if required for the first element.

11. Binary Molecular Compounds N 2 Odinitrogen monoxide N 2 O 3 dinitrogen trioxide N 2 O 5 dinitrogen pentoxide ICliodine monochloride ICl 3 iodine trichloride SO 2 sulfur dioxide SO 3 sulfur trioxide YouTube - Naming molecular compounds

12. Binary Molecular Compounds Containing Two Nonmetals 1.________________ diarsenic trisulfide 2.________________sulfur dioxide 3. P 2 O 5 ____________________ 4.________________ carbon dioxide 5. N 2 O 5 ____________________ 6. H 2 O____________________ As 2 S 3 SO 2 diphosphorus pentoxide CO 2 dinitrogen pentoxide dihydrogen monoxide

13. Naming Binary Compounds Binary Compound? Metal Present? Does the metal form more than one cation? Molecule Use Greek Prefixes Ionic compound (cation has one charge only) Use the element name for the cation. Ionic compound (cation has more than one charge) Determine the Charge of the cation; use a Roman numeral after the cation name. Yes No Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 98

14. Classwork : p 249 #14-18 p 251 #27-29

15. Lesson 3:Molecular Structure Structural formula: uses symbols and bonds to show relative position of atoms. Steps to determine Lewis structures for molecules (p254) 1. Predict the location of certain atoms. Hydrogen is always an end atom The least electronegative atom is the central atom (usually the one closer to the left on periodic table) 2. Find the total number of electrons available for bonding. (# of valence electrons of atoms in molecule) 3. Determine the number of bonding pairs by dividing the total number of electron by 2

16. 4. Place one bonding pair (single bond) between central atom and terminal atoms. 5. Determine number of bonding pairs remaining: Subtract pairs used in step 4 from bonding pairs in step 3. Place lone pairs around each terminal atom bonded to the central atom to satisfy the octet rule. Any remaining pairs are assigned to the central atom. 6. Determine if central atom satisfies the octet rule. If the central atom does not have an octet, convert one or two of the lone pairs on the terminal atoms to a double or a triple bond between central and terminal atom. Some elements do not follow the octet rule: Ex. B, S, P, Xe

17. Ex. 1 Draw the lewis structure for ammonia, NH 3 1. Hydrogen is an end atom and nitrogen is the central atom. 2. Total number of valence electrons: (1 nitrogen x 5 valence electrons)+ (3 hydrogens x 1 valence electron)= 8 valence electrons. 3. Total number of bonding pairs= 8/2 = 4 4. Draw single bond from each H to N N H H H

18. Ex. 1 Draw the lewis structure for ammonia, NH 3 5. Subtract the number of pairs of electrons used from the total pairs of electrons: 4-3 =1 pair available One lone pair remains, hydrogen can have only one bond, assign the lone pair to the central atom, N. N H H H

19. Ex. 2 Draw the lewis structure for carbon dioxide, CO 2 1. Oxygen atoms are end atoms and carbon is the central atom. 2. Total number of valence electrons: (1 carbon x 4 valence electrons)+ (2 oxygen x 6 valence electron)= 16 valence electrons. 3. Total number of bonding pairs= 16/2 = 8 4. Draw single bond from each C to O C OO

20. Ex. 1 Draw the lewis structure for carbon dioxide, CO 2 5. Subtract the number of pairs of electrons used from the total pairs of electrons: 8-2 =6 pair available Add three pairs of electrons to each oxygen. C OO

21. Ex. 1 Draw the lewis structure for carbon dioxide, CO 2 6. No lone pairs remain for carbon. Carbon does not have an octet, use a lone pair from each oxygen to form a double bond with the carbon atom. C OO C OO

22. Learning Check Draw the lewis structure for a)Carbon monoxide, COb)ethylene, C 2 H 4

23. Lewis structures for polyatomic ions The difference is calculating number of valence electrons: First find the number of valence electrons available in the atoms present in the ion. Then subtract the ion charge if the ion is positive and add the ion charge if the ion is negative

24. Ex. Draw the Lewis Structure for the polyatomic ion phosphate, PO 4 3- 1. Find total number of valence electrons: (1 P x 5 valence electrons)+ (4 O x 6 valence electrons) + (3 electrons from negative charge)= 32 2. Determine number bonding pairs: 32/2=16 3. Draw single bond 4. Subtract number of pairs used

25. CW: lewis structures handout part 1

26. Lesson :4 Resonance structures and exceptions to octet rule Resonance: Occurs when more than one valid Lewis structure can be written for a molecule or ion. Differ in the position of electron pairs Ex. O 3, SO 2

27. Exceptions Some molecules have an odd number of valence electrons and cannot form an octet around each atom. Ex. NO 2 Suboctets: Some molecules form with fewer than eight electrons present around an atom. Ex. Boron Expanded Octet: Some compounds have central atoms with more than 8 electrons. This is called an expanded octet. Examples of elements that may have expanded octets: S, P, and Xe.

28. Ex. 3 Draw the lewis structure for XeF 4 (exception octet rule) 1. F is an end atom and nitrogen is the central atom. 2. Total number of valence electrons: (1 xenon x 8 valence electrons)+ (4 fluorines x 7 valence electron)= 36 valence electrons. 3. Total number of bonding pairs= 36/2 = 18 4. Draw single bond from each F to Xe Xe F F F F

29. Ex. 1 Draw the lewis structure for XeF 4 (exception octet rule) 5. Subtract the number of pairs of electrons used from the total pairs of electrons: 18-4 =14 pairs available 14 lone pairs remain, place them around each fluorine so that each fluorine has 8 valence electrons Xe F F F F

30. Ex. 1 Draw the lewis structure for XeF 4 (exception octet rule) 6. There are 2 pairs of electrons still available, place around Xe which is capable of having more than 8 valence electron. Xe F F F F

31. Molecular Shape VSEPR (Valence shell electron pair repulsion) Model The repulsion between electron pairs in a molecule result in atoms existing at fixed angles from each other. (Remember balloon activity) Shared electron pairs repel each other A greater repulsion occurs between unshared electron pairs and shared electron pairs. Read p262 connection to biology.

32. Shape of Molecules: Count number of bonds and unshared pairs of electrons AROUND CENTRAL ATOM and then use table on p263 to determine shape of molecule

33. Use table to determine shape of molecule. Shape: Shared pairs (bonds): Unshared pairs : SO 2 Water molecule: H 2 O 1.Draw Lewis Structure 2.Use table to determine shape

34. Carbon tetrachloride CCl CCl 4 C 109.5 o Cl Shape: Shared pairs (bonds): Unshared pairs:

35. Classwork: p 264 #56-60

36. Lesson 5: Electronegativity and polarity The type of bond can be predicted by using the electronegativity difference (absolute value)of the elements that are bonded.

37. For identical atoms, the bond they form is a nonpolar bond because the pair of electrons is shared equally. Bonds between different atoms can be ionic or covalent. If electronegativity difference is greater than 1.70 it is considered an ionic bond. If electronegativity diffence is 0.4- 1.70 it is considered a polar covalent bond. If electronegativity diffence is < 0.4 it is considered a mostly covalent bond. If electronegativity difference is 0, the bond is nonpolar covalent.

38. Polar covalent bonds Form when pair of electrons is not shared equally by bonding atoms (like a tug-of-war) Partial charges occur at the ends of the bond. Using the symbols  -, partially negative, and  +, partially positive, next to the model of a molecule indicates the polarity of the polar covalent bond.

39. Molecular Polarity Molecules are either polar or nonpolar The nature of the covalent bond and the shape of the molecule result in a polar or nonpolar molecule. Symmetric molecules tend to be nonpolar H-H (H 2 ) has a nonpolar bond thus is a nonpolar molecule. H 2 O has polar bonds and is a polar molecule.

40. CO 2 has polar bonds but do to the molecule’s shape is a nonpolar molecule. Polar bonds in this molecule are opposite to each other and cancel each other, so molecule is nonpolar.

41. CH 4 has polar bonds but do to the molecule’s shape is a nonpolar molecule. Polar bonds in this molecule are opposite to each other and cancel each other, so molecule is nonpolar.

42. CH 3 Cl has polar bonds and do to the molecule’s shape is a polar molecule. Polar bonds in this molecule do not cancel each other, so molecule is polar.

43. Solubility of molecules (ability to dissolve) Polar molecules and ionic compounds are usually soluble in polar substances. Nonpolar molecules dissolve in nonpolar substance. Intermolecular forces (or Van der Waals forces): Weak attraction forces between molecules.

44. CW p 275 #113-122 except 116