1. 4 4-1 © 2006 Thomson Learning, Inc. All rights reserved General, Organic, and Biochemistry, 8e Bettelheim, Brown, Campbell, & Farrell

2. 4 4-2 © 2006 Thomson Learning, Inc. All rights reserved Chapter 4 Chemical Bonds Chemical Bonds

3. 4 4-3 © 2006 Thomson Learning, Inc. All rights reserved Atoms and Their Ions In 1916, Gilbert N. Lewis pointed out that the lack of chemical reactivity of the noble gases indicates a high degree of stability of their electron configurations.

4. 4 4-4 © 2006 Thomson Learning, Inc. All rights reserved The Octet Rule Octet rule: Octet rule: the tendency of group 1A-7A elements to react in ways that achieve an electron configuration of eight valence electrons. cation.An atom that loses one or more electrons becomes a positively charged ion called an cation. anion.An atom that gains one or more electrons becomes a negatively charged ion called a anion.

5. 4 4-5 © 2006 Thomson Learning, Inc. All rights reserved The Octet Rule Example: in losing one electron, a sodium atom forms a sodium ion, which has the same electron configuration as neon. Na (11 electrons): 1s 2 2s 2 2p 6 3s 1 Na + (10 electrons): 1s 2 2s 2 2p 6

6. 4 4-6 © 2006 Thomson Learning, Inc. All rights reserved The Octet Rule Example: in gaining one electron, a chlorine atom forms a chloride ion, which has the same electron configuration as argon. Chlorine atom (17 electrons): 1s 2 2s 2 2p 6 3s 2 3p 5 Chloride ion (18 electrons): 1s 2 2s 2 2p 6 3s 2 3p 6

7. 4 4-7 © 2006 Thomson Learning, Inc. All rights reserved The Octet Rule The octet rule gives us a good way to understand why Group 1A-7A elements form the ions they do; but it is not perfect: Ions of period 1 and 2 elements with charges greater than +2 are unstable. For example, boron does not lose its three valence electrons to become B 3+, nor does carbon lose its four valence electrons to become C 4+ Ions of period 1 and 2 elements with charges greater than -2 are also unstable. For example, carbon does not gain four valence electrons to become C 4- The octet rule does not apply to Group 1B-7B (transition elements), most of which form ions with two or more different positive charges.

8. 4 4-8 © 2006 Thomson Learning, Inc. All rights reserved Naming Cations Elements of Groups 1A, 2A, and 3A form only one type of cation; the name of the cation is the name of the metal followed by the word “ion”.

9. 4 4-9 © 2006 Thomson Learning, Inc. All rights reserved Naming Cations For cations derived from transition and inner transition elements, most of which form more than one type of cation, use either Roman numerals to show charge, or -ous -icuse the suffix -ous to show the lower positive charge, and the suffix -ic to show the higher positive charge.

10. 4 4-10 © 2006 Thomson Learning, Inc. All rights reserved Naming Cations

11. 4 4-11 © 2006 Thomson Learning, Inc. All rights reserved Naming Anions For monatomic (containing only one atom) anions, add “ide” to the stem part of the name. here are the monatomic anions we deal with most often:

12. 4 4-12 © 2006 Thomson Learning, Inc. All rights reserved Polyatomic Ions Table 4.4 Names of Common Polyatomic Ions. Common names, where still widely used, are given in parentheses.

13. 4 4-13 © 2006 Thomson Learning, Inc. All rights reserved Forming Chemical Bonds According to the Lewis model: ionic bondAn atom may lose or gain enough electrons to acquire a filled valence shell and become an ion. An ionic bond is the result of the force of attraction between a cation and an anion. covalentbondAn atom may share electrons with one or more other atoms to acquire a filled valence shell. A covalent bond is the result of the force of attraction between two atoms that share one or more pairs of electrons.

14. 4 4-14 © 2006 Thomson Learning, Inc. All rights reserved Electronegativity Electronegativity: Electronegativity: a measure an atom’s attraction for the electrons it shares in a chemical bond with another atom. on the Pauling scale, fluorine, the most electronegative element is assigned a value of 4.0, and all other elements are assigned values relative to fluorine.

15. 4 4-15 © 2006 Thomson Learning, Inc. All rights reserved Electronegativity Table 4.5 Electronegativity Values of the Elements

16. 4 4-16 © 2006 Thomson Learning, Inc. All rights reserved Ionic Compounds ionic bond According to the Lewis model, an ionic bond is formed by the transfer of one or more valence electrons from an atom of lower electronegativity to an atom of higher electronegativity; anion.the more electronegative atom gains one or more valence electrons and becomes an anion. cation.the less electronegative atom loses electrons and becomes a cation. ionic compound.the compound formed by the combination of an anion and a cation is called an ionic compound.

17. 4 4-17 © 2006 Thomson Learning, Inc. All rights reserved Forming an Ionic Bond In forming sodium chloride, NaCl, one electron is transferred from a sodium atom to a chlorine atom: we use a single-headed curved arrow to show this transfer of one electron:

18. 4 4-18 © 2006 Thomson Learning, Inc. All rights reserved Formulas of Ionic Compounds The total number of positive charges must equal the total number of negative charges. lithium ion and bromide ion form LiBr barium ion and iodide ion form BaI 2 aluminum ion and sulfide ion form Al 2 S 3 sodium ion and bicarbonate ion form NaHCO 3 potassium ion and phosphate ion form K 3 PO 4

19. 4 4-19 © 2006 Thomson Learning, Inc. All rights reserved Naming Ionic Compounds Binary ionic compounds the name of metal from which the positive ion is formed followed by the name of the negative ion; subscripts are ignored: AlCl 3 is aluminum chloride LiBr is lithium bromide Ag 2 S is silver sulfide MgO is magnesium oxide KCl is potassium chloride

20. 4 4-20 © 2006 Thomson Learning, Inc. All rights reserved Naming Ionic Compounds Binary ionic compounds of metals that form two different cations for systematic names, use Roman numerals to show charge on the metal ion; for common names, use the - ous, -ic suffixes CuO is copper(II) oxide; cupric oxide Cu 2 O is copper(I) oxide; cuprous oxide FeO is iron(II) oxide; ferrous oxide Fe 2 O 3 is iron(III) oxide; ferric oxide

21. 4 4-21 © 2006 Thomson Learning, Inc. All rights reserved Naming Ionic Compounds Ionic compounds that contain polyatomic ions name the positive ion first followed by the name of the negative ion. NaNO 3 is sodium nitrate. CaCO 3 is calcium carbonate. NaH 2 PO 4 is sodium dihydrogen phosphate. NH 4 OH is ammonium hydroxide. FeCO 3 is iron(II) carbonate; ferrous carbonate. Fe 2 (CO 3 ) 3 is iron(III) carbonate; ferric carbonate. CuSO 4 is copper(II) sulfate; cupric sulfate.

22. 4 4-22 © 2006 Thomson Learning, Inc. All rights reserved Forming a Covalent Bond A covalent bond is formed by sharing one or more pairs of electrons. the pair of electrons is shared by both atoms and, at the same time, fills the valence shell of each atom. example:example: in forming H 2, each hydrogen contributes one electron to the single bond.

23. 4 4-23 © 2006 Thomson Learning, Inc. All rights reserved Polarity of Covalent Bonds Although all covalent bonds involve sharing of electron pairs, they differ in the degree of sharing: nonpolar covalent bond:nonpolar covalent bond: electrons are shared equally polar covalent bond:polar covalent bond: electron sharing is not equal the degree of sharing depends on the relative electronegativities of the bonded atoms.

24. 4 4-24 © 2006 Thomson Learning, Inc. All rights reserved Polarity of Covalent Bonds Examples:

25. 4 4-25 © 2006 Thomson Learning, Inc. All rights reserved Polarity of Covalent Bonds In a polar covalent bond,  -the more electronegative atom gains a greater fraction of the shared electrons and acquires a partial negative charge; indicated by  - or the head of a crossed arrow  +the less electronegative atom has a lesser fraction of the shared electrons and acquires a partial positive charge; indicated by  + or the tail of a crossed arrow

26. 4 4-26 © 2006 Thomson Learning, Inc. All rights reserved Drawing Lewis Structures 1. Determine the number of valence electrons in the molecule. 2. Determine the arrangement of atoms in the molecule. 3. Connect the atoms by single bonds. 4. Show bonding electrons as a single line; show nonbonding electrons as a pair of Lewis dots. single bond double bond triple bond 5. In a single bond, atoms share one pair of electrons; in a double bond, they share two pairs, and in a triple bond they share three pairs.

27. 4 4-27 © 2006 Thomson Learning, Inc. All rights reserved Lewis Structures Table 4.7 Lewis Structures for Several Small Molecules. Table 4.7 Lewis Structures for Several Small Molecules. (The number of valence electrons is given in parentheses after the molecular formula

28. 4 4-28 © 2006 Thomson Learning, Inc. All rights reserved Lewis Structures Examples draw a Lewis structure for hydrogen peroxide, H 2 O 2. draw a Lewis structure for methanol, CH 3 OH. draw a Lewis structure for acetic acid, CH 3 COOH.

29. 4 4-29 © 2006 Thomson Learning, Inc. All rights reserved Exceptions to the Octet Rule Atoms of period 2 elements use 2s and 2p orbitals for bonding: these four orbitals can contain a maximum of 8 electrons; hence the octet rule. Atoms of period 3 elements have one 3s orbital, three 3p orbitals, and five 3d orbitals: these nine orbitals can accommodate more than eight electrons, by using 3d orbitals; period 3 atoms can have more than eight electrons in their valence shells.

30. 4 4-30 © 2006 Thomson Learning, Inc. All rights reserved Exceptions to the Octet Rule Phosphorus Phosphorus Sulfur Sulfur

31. 4 4-31 © 2006 Thomson Learning, Inc. All rights reserved Molecular Compounds Molecular compound: Molecular compound: a compound in which all bonds are covalent. Naming binary molecular compounds: the less electronegative element is named first (it is generally written first in the formula). prefixes “di-”, tri-”, etc. are used to show the number of atoms of each element; the prefix “mono-” is omitted when it refers to the first atom, and is rarely used with the second atom. Exception: carbon monoxide NO is nitrogen oxide (nitric oxide) SF 2 is sulfur difluoride N 2 O is dinitrogen oxide (laughing gas)

32. 4 4-32 © 2006 Thomson Learning, Inc. All rights reserved VSEPR Model Valence-Shell Electron-Pair Repulsion (VSEPR) Model valence electrons of an atom may be involved in forming single, double, or triple bonds, or they may be unshared. each combination creates a negatively charged region of electron density around the nucleus. because like charges repel each other, the various regions of electron density around an atom spread so that each is as far away from the others as possible.

33. 4 4-33 © 2006 Thomson Learning, Inc. All rights reserved VSEPR Model Predict the shape of methane, CH 4 The Lewis structure shows carbon surrounded by four regions of electron density. According to the VSEPR model, the four regions radiate from carbon at angles of 109.5°, and the shape of the molecule is tetrahedral. The measured H-C-H bond angles are 109.5°.

34. 4 4-34 © 2006 Thomson Learning, Inc. All rights reserved VSEPR Model Predict the shape of ammonia, NH 3 The Lewis structure shows nitrogen surrounded by four regions of electron density; three regions contain single pairs of electrons, and the fourth contains an unshared pair of electrons. According to the VSEPR model, the four regions radiate from nitrogen at angles of 109.5°, and the shape of the molecule is pyramidal. The measured H-N-H bond angles are 107.3° The unshared pair is not shown on this model.

35. 4 4-35 © 2006 Thomson Learning, Inc. All rights reserved VSEPR Model Predict the shape of water, H 2 O The Lewis structure shows oxygen surrounded by four regions of electron density; two regions contain single pairs of electrons, and the third and fourth contain unshared pairs of electrons. According to the VSEPR model, the four regions radiate from oxygen at angles of 109.5°, and the shape of the molecule is bent. The measured H-O-H bond angle is 104.5°. The unshared pairs are not shown on the model.

36. 4 4-36 © 2006 Thomson Learning, Inc. All rights reserved VSEPR Model Predict the shape of formaldehyde, CH 2 O The Lewis structure shows carbon surrounded by three regions of electron density; two regions contain single pairs of electrons and the third contains two pairs of electrons forming the double bond to oxygen. According to the VSEPR model, the three regions radiate from carbon at angles of 120°, and the shape of the molecule is planar (trigonal planar). The measured H-C-H bond angle is 116.5°.

37. 4 4-37 © 2006 Thomson Learning, Inc. All rights reserved VSEPR Model Predict the shape of ethylene, C 2 H 4 The Lewis structure shows carbon surrounded by three regions of electron density; two regions contain single pairs of electrons and the third contains two pairs of electrons forming the double bond to the other carbon. According to the VSEPR model, the three regions radiate from carbon at angles of 120°, and the shape of the molecule is planar (trigonal planar). The measured H-C-H bond angle is 117.2°.

38. 4 4-38 © 2006 Thomson Learning, Inc. All rights reserved VSEPR Model Predict the shape of acetylene, C 2 H 2 The Lewis structure shows carbon surrounded by two regions of electron density; one region contains a single pair of electron, and the second contains three pairs of electrons forming the triple bond to carbon. According to the VSEPR model, the two regions radiate from carbon at an angle of 180°, and the shape of the molecule is linear. The measured H-C-C bond angle is 180°.

39. 4 4-39 © 2006 Thomson Learning, Inc. All rights reserved

40. 4 4-40 © 2006 Thomson Learning, Inc. All rights reserved Resonance For many molecules and ions, no single Lewis structure provides a truly accurate representation:

41. 4 4-41 © 2006 Thomson Learning, Inc. All rights reserved Resonance Linus Pauling - 1930s Many molecules and ions are best described by writing two or more Lewis structures. contributing structuresIndividual Lewis structures are called contributing structures. Connect individual contributing structures by double- headed (resonance) arrows. hybridThe molecule or ion is a hybrid of the various contributing structures.

42. 4 4-42 © 2006 Thomson Learning, Inc. All rights reserved Resonance Examples

43. 4 4-43 © 2006 Thomson Learning, Inc. All rights reserved Resonance Curved arrow: Curved arrow: a symbol used to show the redistribution of valence electrons. In using curved arrows, there are only two allowed types of electron redistribution: from a bond to an adjacent atom. from an atom to an adjacent bond.

44. 4 4-44 © 2006 Thomson Learning, Inc. All rights reserved Resonance All contributing structures must 1. have the same number of valence electrons. 2. obey the rules of covalent bonding. no more than 2 electrons in the valence shell of H. no more than 8 electrons in the valence shell of a 2nd period element. 3rd period elements, such as P and S, may have up to 12 electrons in their valence shells. 3. differ only in distribution of valence electrons; the position of all nuclei must be the same. 4. have the same number of paired and unpaired electrons.

45. 4 4-45 © 2006 Thomson Learning, Inc. All rights reserved Polarity of Molecules A molecule will be polar if: it has polar bonds, and its centers of partial positive and partial negative charges lie at different places within the molecule. Carbon dioxide, CO 2, has two polar bonds but, because of its geometry, is a nonpolar molecule.

46. 4 4-46 © 2006 Thomson Learning, Inc. All rights reserved Polarity of Molecules Water, H 2 O, has two polar bonds and, because of its geometry, is a polar molecule.

47. 4 4-47 © 2006 Thomson Learning, Inc. All rights reserved Polarity of Molecules Ammonia, NH 3, has three polar bonds and, because of its geometry, is a polar molecule.

48. 4 4-48 © 2006 Thomson Learning, Inc. All rights reserved Polarity of Molecules Both dichloromethane, CH 2 Cl 2, and formaldehyde, CH 2 O, have polar bonds and are polar molecules.

49. 4 4-49 © 2006 Thomson Learning, Inc. All rights reserved End Chapter 4 Chemical Bonds