Larry Emme Chemeketa Community College

1. Chemical Bonds Chapter 11 Larry Emme Chemeketa Community College

2. Periodic Trends in Atomic Properties

3. Characteristic properties and trends of the elements are the basis of the periodic table’s design.

4. These trends allow us to use the periodic table to accurately predict properties and reactions of a wide variety of substances.

5. Metals and Nonmetals

6. metals tend to lose electrons and form positive ions (cations). nonmetals tend to gain electrons and form negative ions (anions). Chemical Properties of Metals Chemical Properties of Nonmetals When metals react with nonmetals electrons are usually transferred from the metal to the nonmetal.

7. Metals are found to the left of the metalloids Nonmetals are found to the right of the metalloids. 11.1

8. Atomic Radius

9. Radii of atoms increase down a group. For each step down a group, electrons enter the next higher energy level.

10. Radii of atoms tend to decrease from left to right across a period. For representative elements within the same period the energy level remains constant as electrons are added. Each time an electron is added a proton is a added to the nucleus. This increase in positive nuclear charge pulls all electrons closer to the nucleus.

11. Ionization Energy

12. The ionization energy of an atom is the energy required to remove the outermost electron from an atom. Na + ionization energy → Na+ + e-

13. Ionization energies gradually increase from left to right across a period. Noble Gases 1 2 Periodic relationship of the first ionization energy for representative elements in the first four periods. VIIA VA 3 VIA IA 4 IVA IIA IIIA

14. nonmetals metals Ionization energies of Group A elements decrease from top to bottom in a group. Noble Gases nonmetals have higher ionization potentials than metals Periodic relationship of the first ionization energy for representative elements in the first four periods. VIIA VA VIA IA IVA Distance of Outer Shell Electrons From Nucleus IIA IIIA

15. Lewis “Dot” Structures of Atoms

16. Metals form cations and nonmetals form anions to attain a stable valence electron structure.

17. These rearrangements occur by losing, gaining, or sharing electrons. This stable structure often consists of two s and six p electrons.

18. The Lewis structure of an atom is a representation that shows the valenceelectrons for that atom. • Na with the electron structure 1s22s22p63s1 has 1 valence electron. • Fluorine with the electron structure 1s22s22p5 has 7 valence electrons

19. The Lewis structure of an atom uses dots to show the valence electrons of atoms. Unpaired electron B Paired electrons Symbol of the element 2s22p1 The number of dots equals the number of s and p electrons in the atom’s outermost shell.

20. The Lewis structure of an atom uses dots to show the valence electrons of atoms. S 3s23p4 The number of dots equals the number of s and p electrons in the atom’s outermost shell.

21. Lewis Structures of the first 20 elements.

22. The chemistry of many elements, especially the representative ones, is to attain the same outer electron structure as one of the noble gases.

23. With the exception of helium, this structure consists of eight electrons in the outermost energy level.

24. The Covalent Bond:Sharing Electrons

25. A covalent bond consists of a pair of electrons shared between two atoms. In the millions of chemical compounds that exist, the covalent bond is the predominant chemical bond.

26. Substances which covalently bond exist as molecules. Carbon dioxide bonds covalently. It exists as individually bonded covalent molecules containing one carbon and two oxygen atoms.

27. The term molecule is not used when referring to ionic substances. Sodium chloride bonds ionically. It consists of a large aggregate of positive and negative ions. No molecules of NaCl exist.

28. The most likely region to find the two electrons is between the two nuclei. The orbital of the electrons includes both hydrogen nuclei. Two 1s orbitals from each of two hydrogen atoms overlap. Two 1s orbitals from each of two hydrogen atoms overlap. Each 1s orbital contains 1 electron. The two nuclei are shielded from each other by the electron pair. This allows the two nuclei to draw close together. Covalent bonding in the hydrogen molecule

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

30. Electronegativity Linus Pauling

31. electronegativity The relative attraction that an atom has for a pair of shared electrons in a covalent bond.

32. If the two atoms that constitute a covalent bond are identical then there is equal sharing of electrons. • This is called nonpolar covalent bonding. • Ionic bonding and nonpolar covalent bonding represent two extremes.

33. If the two atoms that constitute a covalent bond are not identical then there is unequal sharing of electrons. • This is called polar covalent bonding. • One atom assumes a partial positive charge and the other atom assumes a partial negative charge. • This charge difference is a result of the unequal attractions the atoms have for their shared electron pair.

34. Polar and Non-Polar

35. + - H Cl Chlorine has a greater attraction for the shared electron pair than hydrogen. Partial positive charge on hydrogen. Partial negative charge on chlorine. Polar Covalent Bonding in HCl : : The attractive force that an atom of an element has for shared electrons in a molecule or a polyatomic ion is known as its electronegativity. Shared electron pair. The shared electron pair is closer to chlorine than to hydrogen.

36. A scale of relative electronegativities was developed by Linus Pauling.

37. Electronegativity decreases down a group for representative elements. Electronegativity generally increases left to right across a period.

38. + - A dipole can be written as A dipole is a molecule that is electrically asymmetrical, causing it to be oppositely charged at two points.

39. O H Br H Cl H H An arrow can be used to indicate a dipole. The arrow points to the negative end of the dipole. Molecules of HCl, HBr and H2O are polar .

40. A molecule containing different kinds of atoms may or may not be polar depending on its shape. The carbon dioxide molecule is nonpolar because its carbon-oxygen dipoles cancel each other by acting in opposite directions.

41. Lewis Structures ofCompounds

42. In writing Lewis structures, the most important consideration for forming a stable compound is that the atoms attain a noble gas configuration.

43. The most difficult part of writing Lewis structures is determining the arrangement of the atoms in a molecule or an ion. • In simple molecules with more than two atoms, one atom will be the central atom surrounded by the other atoms.

44. Cl2O has two possible arrangements. The two chlorines can be bonded to each other. Cl-Cl-O The two chlorines can be bonded to oxygen. Cl-O-Cl Usually the single atom will be the central atom.

45. Practice WritingLewis Structures

46. Valence Electrons of Group A Elements

47. 3-Dimensional Shapes Linear 180 Bent 105 Trigonal Planar 120 Tetrahedral 109.5 Trigonal Pyramidal 107

48. Covalent Bonding Structures Bent 105 Polar 180 Linear Non-Polar Trigonal Pyramidal 107 Polar Trigonal Planar 120 Non-Polar Tetrahedral Non-Polar 109.5

49. “Define ‘resonance’? Sure, that’s where you live.”

50. The Ionic Bond: Transfer ofElectrons From One Atomto Another