Chapter 6.1

1. Chapter 6.1 Introduction to Chemical Bonding

2. Why do elements bond? • They want to become more stable elements, which they achieve by having 8 valence electrons, causing a decrease in the atom’s potential energy.

3. Types of Chemical Bonding • When atoms bond, their valence electrons are moved around in a way to make the atom more stable. • Ionic Bonding: the electrical attraction between cations and anions. • Covalent Bonding: the sharing of electrons between non-metals

4. Ionic or Covalent? • Looking at the difference in electronegativity of two elements will tell us which bonding it will favor. • Electronegativity difference of 1.7 or less will favor covalent and a difference of 1.8 or more will favor ionic.

5. Types of Covalent • Nonpolar-covalent: electrons are shared equally. Difference of 0.0 – 0.3 • Polar-Covalent: electrons are not shared equally, one atom will have a stronger pull on the electrons. Difference of 0.4 to 1.7

6. Example • Use electronegativity differences to classify bonding between sulfur, and hydrogen, cesium, and chlorine. In each pair, which atom will be more negative. • Look at figure 20 on page 161 to find electronegativities.

7. Sulfur and Hydrogen • 2.5-2.1 = 0.4 • polar-covalent • sulfur • Sulfur and Cesium • 2.5-0.7 = 1.8 • ionic • sulfur • Sulfur and Chlorine • 3.0-2.5 = 0.5 • polar-covalent • chlorine

8. Chapter 6.2 Covalent Bonding and Molecular Compounds Water Fructose Carbon Dioxide Ammonia

9. Why Do Atoms Bond? To get eight valence electrons To become more stable In ionic bonds, metals lose electrons and non-metals gain electrons. What happens when both elements need electrons?

10. Molecules and Molecular Compounds Compounds that are NOT held together by an electrical attraction, but instead by a sharing of electrons. Atoms held together by sharing electrons and filling the outer energy levels are joined by a covalent bond. NONMETALS ONLY!! - No metals

11. Molecules and Molecular Compounds A molecule is a neutral group of atoms joined together covalent bonds. A compound composed of molecules is called a molecular compound. The chemical formula for a molecule is called the molecular formula. A chemical formula tells you how many atoms of each element one molecule of a compound contains.

12. Learning Check Indicate whether a bond between the following would be 1) Ionic 2) covalent ____ A. sodium and oxygen ____ B. nitrogen and oxygen ____ C. phosphorus and chlorine ____ D. calcium and sulfur ____ E. chlorine and bromine Ionic Covalent Covalent Ionic Covalent

13. Monatomic (One Atom) • Noble gases are monatomic. • They exist as single atoms and do not combine with any other elements. • Ex: He, Ne, Ar, Kr, Xe, Rn

14. 7 Diatomic Molecules Some elements will covalently bond to themselves to form a molecule composed of TWO atoms. Some elements occur as “diatomic” molecules in nature because they are more stable than individual atoms The 7 diatomic elements are all gases: H2, O2, N2, Cl2, Br2, I2, F2

15. Strength of Covalent Bonds Distance between two bonding nuclei at the position of maximum attracting is bond length Bond length is determined by the size of the atoms and how many electron pairs are shared Bond energy is the energy required to break a chemical bond and form neutral isolated atoms.

16. Homework • Page 177 #1-3 • Page 209 #6, 10-12

17. 6.2 Notes Continued

18. Octet Rule in Covalent Bonds • Remember that all compounds want to attain the electron configuration of noble gases. • Hydrogen only needs 2, the rest need 8. • Regarding covalent bonds, electrons are shared between the atoms so that they attain the electron configuration of noble gases.

19. H - B - H H N O O Exceptions to the Octet Rule • A small group of molecules has an odd number of valence electrons and cannot form an octet around each atom -Ex: NO2 • Fewer than eight electrons: BORON is stable with 6! • Ex: BH3

20. Exceptions to the Octet Rule • Some central atoms have more than eight valence electrons • Referred to as an “expanded” octet • Explained by d-orbitals PCl5 (10 e-) SF6 (12 e-)

21. Drawing Valence Electrons • “Electron-dot notation”: Electrons are represented as dots located around the symbol of the element. You must put one electron on each side before you double up. Examples: Nitrogen = Hydrogen = Carbon = X N H C

22. In-Class Examples • Chlorine • Neon • Magnesium • Sulfur • Silicon

23. To draw Lewis structures for covalent bonds, use the NASB method: N (Needed): Find the number of electrons needed to form full octets for all elements. For most nonmetals, they need 8. Hydrogen needs only 2. A (Available): Find the number of electrons available by adding up all of the valence electrons for all elements involved. S (Shared): Subtract the two numbers. S= N-A B (Bond): A bond is formed with 2 electrons, so divide by 2 to find how many bonds to draw between the elements. Draw the molecule. Put first atom in the center. H’s are always outside. Draw in the bonds, then fill in the rest of the electrons. Check to ensure all atoms have a full octet.

24. Draw the Lewis-dot-structure for the following molecules 1. HF 2. CCl2H2

25. Draw the Lewis-dot-structure for the following molecules 1. H2O 2. CO2

26. Types of Bonds Each bond involves the sharing of _____ _________ of electrons. Single Bonds= __ e-’s Double Bonds= __ e-’s Triple Bonds=__ e-’s one pair 2 4 6

27. Resonance Structures • Occurs when more than one valid Lewis Structure can be written for a molecule or ion • Differ only in the position of electron pairs, never the atom’s positions • Actual molecule behaves as if it has one structure • Example: O3

28. Homework • 6.2 page 209 #15-19, 21, 23

30. Bond formed between two or more ions to form an electrically neutral compound by the transfer of electrons. Formula Unit: the simplest collection of atoms from which an ionic compound’s formula can be established. Ch 6.3 IONIC BONDS

31. Ionic Bonding How Ionic Bonding Works The negative and positively-charged ions are attracted to each other (like a magnet). ****Ionic bonding – only 2types 1 Metal ion + 1 Nonmetal ion or 1 Metal ion + 1 Polyatomic ion

33. Ionic Bonds: One Big Greedy Thief Dog!

34. Ionic bond – electron from Na is transferred to Cl.

35. Characteristics of Ionic and Covalent Compounds

36. Polyatomic Ions • A charged group of covalently bonded atoms. • They behave as one group. • If more than one is needed, written with parentheses around the ion. • NaCl2 vs. Mg(OH)2 • They are held together by covalent bonds, but form ionic bonds with other ions.

37. Ch 6.4 Bonding in Metals Metallic bonds -Bonds found in metals -Holds metal atoms together very strongly.

38. Metallic Bond • Good conductors at all states, shiny, very high melting points • The valence electrons of metal atoms can be modeled as a sea of electrons and the electrons can move freely. • Malleability and Ductile

39. Crystalline Structure of Metals • Metal atoms are arranged in very compact and orderly patterns. • Resembles how apples and oranges are stacked in a grocery store.

40. Homework • 6.3 and 6.4 pg 210 #25-26, 28, 30-31

41. Ch 6.5 Molecular Geometry

42. VSEPR Theory • Use VSEPR theory • Valence-Shell-Electron-PairRepulsion • =the repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible. • Determined by number of valence electrons of the central atom • 3-D shape is a result of bonded pairs and lone pairs of electrons

43. Bonding and Shape of Molecules Number of Bonds Number of Unshared Pairs Covalent Structure Shape Examples AB2 0 0 0 1 1 2 3 4 3 2 Linear Trigonal planar Tetrahedral Pyramidal Bent or Angular BeCl2 BF3 CH4, SiCl4 NH3, PCl3 ONF AB3 AB4 AB3E AB2E

45. .. .. .. S O O O C O O S N C F F O O F F F F F F F F F P S F F F F F F The VSEPR Model SO2 Linear Bent Trigonal planar Trigonal pyramidal AB6 F Tetrahedral Trigonal bipyramidal Octahedral

46. H H C H H C H H H 109.5o H Tetrahedral geometry Methane CH4

47. .. .. N H H H N H H 107o H N H H H Ammonia- NH3 N H H H Trigonal Pyramidal geometry

48. .. .. .. .. O H H .. .. Water –H2O SO2 O H H Bent geometry

49. F F B B F F F F 120o Boron trifluoride - BF3 Trigonal planar

50. C O O Carbon dioxide – CO2 O C O Linear