1. Bonding Forces of attraction that hold atoms together making compounds

2. Chemical symbols Symbols are used to represent elements Either one capital letter, or a capital letter with a lower case letter Know names and symbols of elements: 1 – 30, plus Rb, Cs, Sr, Ba, Ag, Au, Cd, Hg, Pt, Ga, Ge, As, Sn, Pb, Se, Br, I, and U

3. Basic idea... All chemical bonds form because they impart stability to the atoms involved lower energy = greater stability

4. Quick review All types of chemical bonds involve electrons Valence electrons, the electrons in the outermost occupied energy level of an atom, are usually the electrons involved in bonding

5. The representative elements have the same number of valence electrons as their family number in the American system Example: Mg, column IIA, 2 valence electrons The transition metals all have two valence electrons  ns 2 (n-1)d x

6. Lewis dot structures are used to represent the valence electrons each dot represents a valence electron no more than 8 dots total no more than 2 dots on a side.

7. Lewis dot structures of representative elements

8. The Octet Rule Atoms will gain, lose, or share electrons in order to achieve an ns 2 np 6 valence configuration – 8 valence electrons

9. cations Atoms that have lost one or more electrons Become positive (+) ions Usually the metals Only have 1,2 or 3 valence electrons

10. anions Atoms that have gained one or more electrons Become negative (-) ions Usually nonmetals Have 5, 6, or 7 valence electrons Gain enough electrons to end up with 8 valence electrons

11. Yellow = metals Green = semimetallics Blue + purple = nonmetals

12. Ionic Bonding Metals lose electrons easily, nonmetals have a strong attraction for more electrons metal atoms will lose electrons to nonmetal atoms, causing both to become ions

13. Metals, having lost one or more electrons, become cations (+) Nonmetals, having gained one or more electrons, become anions (-) Opposites attract: the cations and anions are held together electrostaticly termed “ionic bonds”

14. In summary... Ionic bonds are electrostatic attractions between cations and anions that are formed when electron(s) are transferred from the metal to the nonmetal.

15. Ionic Compounds 1. Only exist as compounds as solids 2. Held together by ionic bonds - A strong attractive force 3. High melting points - Must break the bonds to melt the solid - The higher the melting point, the stronger the ionic bonds

16. 4. brittle solids 5. nonconducting as solids - No charges can flow 6. conduct electricity as liquids or aqueous - Ions are free to move Ionic Compounds

17. As solids, exist in a 3-D repeating pattern called a crystal “lattice” the lattice energy is the energy lowering (stability) accomplished by the formation from “free” ions Also a measure of the energy required to break apart the ionic compound once formed The greater the lattice energy, the stronger the force of attraction

18. Cation (+) Ionic compound = crystalline solid

19. Ion dissociation Many ionic compounds will dissolve in water if it results in more stability (lower E) than in the solid ionic compound the ions “dissociate” from each other Ex: CaCl 2(s) + H 2 O  Ca 2+ (aq) + 2Cl - (aq)

20. Ionic Bond Strength A measure of the attractive force between the ions smaller atoms = stronger ionic bonds fewer atom ratio = stronger bond evidence: melting points

21. Compare the melting points: KCl : 776 o C KI : 723 o C smaller atoms result in stronger ionic bonds

22. Compare the melting points: CaCl 2 : 772 o C NaCl : 800 o C fewer atoms result in stronger ionic bonds

23. Periodic trends…

24. Sizes of atoms Periodic trend: atomic radii increase moving down a group Increasing energy level Periodic trend: atomic radii decrease moving left to right in a period The charge felt by the valence electrons becomes larger

25. There is a general decrease in atomic radius from left to right, caused by increasing positive charge in the nucleus. Valence electrons are not shielded from the increasing nuclear charge because no additional electrons come between the nucleus and the valence electrons. Sizes of atoms

26. For elements that occur as molecules, the atomic radius is half the distance between nuclei of identical atoms. For metals, atomic radius is half the distance between adjacent nuclei in a crystal of the element.

27. Atomic Radius

28. Atomic radius generally increases as you move down a group. The outermost orbital size increases down a group, making the atom larger.

29. Ionization energy The energy needed to remove a valence electron from an atom A measure of how tightly the electrons are being held periodic trend increases from the bottom up increases left to right

30. In general, metals have lower IE than nonmetals alkali metals are the lowest IE family noble gases are highest IE family

31. The energy required to remove the first electron is called the first ionization energy. Ionization energy First ionization energy increases from left to right across a period. First ionization energy decreases down a group because atomic size increases and less energy is required to remove an electron farther from the nucleus.

32. Ionization energy

34. Electron affinity A measure of how strongly an element would like to gain an electron periodic trend increases from the bottom up increases left to right ignore the noble gases

35. Atoms that lose electrons easily have little attraction for additional electrons (and vice versa) metals have low IE, low EA Nonmetals have high IE, high EA Octet rule: when atoms react, they tend to strive to achieve a configuration having 8 valence electrons This results in some form of bond formation

36. Periodic trends… As you move from left to right along a period… Atoms get …. Smaller Ionization energy goes …. Up Electron affinity goes …. Up

37. Periodic trends… As you move down a group/family Atoms get …. Larger Ionization energy goes …. Down Electron affinity goes …. Down

38. Check your understanding The lowest ionization energy is the ____. A.first B.second C.third D.fourth

39. Bonding Forces of attraction that hold atoms together making compounds

40. Types of compounds All compounds are made of two or more elements held together by chemical bonds Ions of opposite charges are held together by ionic bonds Ionic bonding is non-directional There are no “ionic molecules” Formulas of ionic compounds show the ratio of cation to anion Ionic compounds only exist in the solid state, in a 3-D crystal lattice

41. Covalent Bonding Covalent bonding involves the sharing of electron pairs usually between two high EA, high IE nonmetals both want more e - ’s, neither is willing to lose the e - ’s they have

42. A nonmetal will form as many covalent bonds as necessary to fulfill the octet rule example: C, with 4 valence e - ’s, will form 4 covalent bonds results in 8 valence e - ’s around the carbon atom at least part of the time double and triple covalent bonding is a possibility Covalent Bonding

43. Covalent compounds are made of two or more elements held together by covalent bonds Covalent bonding is directional Between two individual atoms A group of covalently bonded atoms is referred to as a “molecule” Covalent compounds are also referred to as “molecular” compounds Types of compounds

44. When does the octet rule fail?

45. H, He and Li Helium strives for 2 valence electrons 1s 2 configuration Hydrogen will sometimes will share its one electron with another atom, forming a single covalent bond Lithium will lose its lone valence electron, gaining the 1s 2 configuration of He

46. Be Be will sometimes lose its 2 valence electrons, gaining the Is 2 configuration of He Be will sometimes form 2 covalent bonds, giving it 4 valence electrons nuclear charge of +4 cannot handle 8 valence electrons

47. B Boron will often make three covalent bonds using its three valence electrons nuclear charge of +5 cannot handle 8 valence electrons in a stable manner

48. “organometallic” compounds Some metals will form covalent compounds with nonmetals Hg, Ga, Sn, and others The octet rule is not followed for the metals,but is for nonmetals Form 2 or more covalent bonds

49. P, S, Cl, Se, Br, I Elements in the third period and lower have empty d orbitals there is room for more than 8 valence electrons These elements will at times make more than 4 covalent bonds

50. Rules for Drawing structural formulas 1) Determine the central atom, place the other atoms evenly spaced around the outside 2) Count the total number of valence electrons 3) Draw single bonds between the central atoms and each of the outside atoms

51. 4) Complete the octet on the outside atoms by placing electrons in pairs around the outside atoms (lone pairs) 5) Place any remaining electrons on the central atom in pairs 6) If the central atom does not have its minimum number of electrons (usually 8), form double bonds by moving lone pairs off of the outside atoms and drawing them as bonding pairs **never make double bonds until ALL of the electrons are used first!!