Periodic properties of the elements

1. Periodic properties of the elements AP Chemistry Chapter 7

2. Development of the periodic table • 1869: Dmitri Mendeleev • Arranged elements by increasing atomic mass • Grouped elements with similar properties in columns • Left blank spaces for undiscovered elements • Predicted the properties of Ga & Ge before their discoveries • 1913: Henry Moseley • Developed concept of atomic numbers (after Rutherford’s Nuclear Model) • Identified this as the # of protons in the nucleus of the atom • Arranged the periodic table in based on atomic #’s • All elements grouped with similar properties (Ex. Ar & K) • Allowed for “holes” for undiscovered elements

3. Effective nuclear charge • Force of attraction between the nucleus and the electrons • Coulomb’s law: the strength of the interaction between 2 electrical charges (attraction/repulsion) depends on the magnitudes of the charges and the distance between them • Farther away = less force Greater magnitude of charge = more force • Each electron is simultaneously attracted by the nucleus and repelled by other electrons • We can only estimate the net attraction of each electron by the nucleus by considering how it interacts with the average environment created by the nucleus and other electrons in the atom. • We view this net electric field as if it results from a single positive charge located at the nucleus…..We call this the effective nuclear charge (Zeff)

4. Effective nuclear charge • Zeff is smaller than the actual charge (Z) because it takes into account the repulsion between the electrons • A valence electron is attracted to the nucleus and repelled by other electrons in the atom • The inner electrons partially cancel (shield) the attraction between the nucleus and the valence electrons Screening constant # of protons Effective nuclear charge Screening constant is usually the # of core electrons in the atom

5. Effective nuclear charge • Example: Na • This is only a rough estimate…the situation is more complicated because of electron distributions of atomic orbitals (uncertainty)

6. Periodic trend for zeff • Zeff tends to increases for valence electrons from left to right across a period on the periodic table • Ex: Li & Be • As you move down a group (column), the Zeff changes far less than it does as you move across a row. (It does increase slightly – but it is less important than the increase in moving across a period)

7. Sizes of atoms & ions • According to the quantum mechanical model, atoms and ions do not have sharply defined boundaries where electron distribution becomes zero….leads to different definitions of atomic size • Nonbonding radius of atoms: The closest distance between colliding atoms without forming bonds • Bonding atomic radius: The attractive interaction between two atoms brings them closer together than in the non-bonding situation • Bonding atomic radius is shorter than the nonbonding radius

9. Sizes of atoms & ions • Space filling models are based on nonbonding radius (van der Waals radius) • Space filling models of molecules uses bonding atomic radii (covalent radii) to determine the amount of overlap of the atoms • Bonding atomic radii can be assigned based on measurements of distances between nuclei in molecules • Ex. I2 distance between I – I is 2.66 angstroms assign a bond radius of 1.33 angstroms for I Cl2 C – C C – Cl

11. Sample exercise 7.1 • Natural gas used in home heating and cooking is odorless. Because natural gas leaks pose the danger of explosion or suffocation, various smelly substances are added to the gas to allow detection of a leak. One such substance is methyl mercaptan, CH3SH, whose structure is shown. Use Figure 7.6 to predict the lengths of the C – S, C – H, and S – H bonds in this molecule.

12. Practice exercise page 267 • Using Figure 7.6, predict which will be greater, the P – Br bond length in PBr3 or the As – Cl bond length in AsCl3

13. Periodic trends in atomic radii • Increases within a group (top to bottom) • Decreases left to right across a period • These trends apply to representative elements • The increase down a group has a larger effect than the decrease across a period

14. Sample exercise 7.2 • Referring to a periodic table, arrange (as much as possible) the following atoms in order of increasing size: P, S, As, Se

15. Practice exercise page 268 • Arrange the following atoms in order of increasing atomic radius: Na, Be, Mg

16. Periodic trends in ionic radii • Radii of ions are based on the distances between nuclei in ionic compounds • Formation of a cation vacates the outer most occupied orbital of the atom and decreases electron – electron repulsion • Results is that cations are smaller than their parent atoms • When electrons are added in the formation of anions Electron – electron repulsions cause the electrons to spread out more in space Result is that anions are larger than their parent atoms

17. Periodic trends in atomic radii • For ions with the same charge, size increases as we move down a column in the periodic table

18. Sample Exercise 7.3 • Arrange these atoms and ions in order of decreasing size: Mg2+ , Ca2+, and Ca.

19. Practice Exercise (page 268) • Which of the following atoms and ions is largest: S2- , S, O2-

20. Isoelectronic series • A group of ions all containing the same number of electrons is an isoelectronic series • When isoelectronic series is arranged by increasing atomic # (increasing nuclear charge)….radius decreases (larger nuclear charge on the same # of electrons) • Example arrange: Na+, O2-, Al3+, F-, Mg2+

21. Sample Exercise 7.4 • Arrange the ions K+, Cl-, Ca2+, and S2- in order of decreasing size.

22. Practice Exercise (page 270) • Which of the following is largest, Rb+, Sr2+ , or Y3+ ?

23. Ionization Energy • Ionization energy is the minimum amount of energy required to remove an electron from an atom or ion • First ionization energy is energy required to remove an electron from ground state of the gaseous atom • Na(g)  Na+ (g) + e- (1st ionization energy) • Na+ (g)  Na2+ (g) + e- (2nd ionization energy) • Greater ionization energy = more difficulty in removing an electron

24. Successive Ionization Energies • Table 7.2 • * note sharp increase in ionization energy to remove inner shell electrons

26. Sample Exercise 7.5 • Three elements are indicated in the periodic table in the margin. Based on their locations, predict the one with the largest second ionization energy. (Elements are sulfur, sodium, and calcium)

27. Practice Exercise • Which will have the greater third ionization energy, Ca or S?

28. Periodic Trends in First Ionization Energy • Figure 7.10 • Within a period, I1 increases left  right (alkali metals lowest  noble gases highest) • Within a column (group) I1 decreases top  bottom • Representative elements show a larger variation in I1 than transition metals do

29. Some exceptions • B has lower ionization energy than Be • O has lower ionization energy than N

30. Sample Exercise 7.6 • Referring to a periodic table, arrange the following atoms in order of increasing first ionization energy: Ne, Na, P, Ar, K

31. Practice Exercise (page 274) • Which has the lowest first ionization energy, B, Al, C, or Si? Which has the highest first ionization energy?

32. Electron Configurations of Ions • When electrons are removed from atoms to form cations, they are always removed from the orbitals with the largest principal quantum numbers first (4s before 3d) • Li  • Fe2+ • Fe3+

33. Sample Exercise 7.7 • Write the electron configuration for (a) Ca2+ (b) Co3+ (c) S2-

34. Practice Exercise (page 275) • Write the electron configuration for: (a) Ga3+ (b) Cr3+ (c) Br-

35. Electron Affinities • Electron affinity is the energy change that occurs when an electron is added to a gaseous atom • Ionization energy: • Cl(g)  Cl-(g) + e- ∆E = 1251 kJ/mol requires input of energy • Cl(g) + e-  Cl-(g) ∆E = -349 kJ/mol releases energy More negative electron affinity = a greater attraction between the atom and the added electron

36. Electron Affinity If ∆E > 0 (like in noble gases)….means that the ion is unstable and will not form

37. Trends in Electron Affinity • Halogens: • Have most negative electron affinities • Gaining one electron gives them a noble gas configuration • Noble Gases: • Have highly positive electron affinity • Group 5A (15): • Have ½ filled p subshells • An additional electron would go into an already occupied orbital resulting in greater electron – electron repulsion • Electron affinity is less negative than elements to the left • Not a large difference in electron affinity as you move down in a group due to the larger orbital (more spread out results in less electron – electron repulsion) and the electron – nucleus attraction also decreases

38. Metals • Middle & left of periodic table • Metallic character generally decreases left to right and increases down a group • Conduct heat & electricity • Malleable • Ductile • Shiny luster • Solids at room temp (except Hg) • Low ionization energy • Most metal oxides are basic: • Na2O(s) + H2O(l)  2NaOH(aq)

39. Sample Exercise 7.8 • (a) Would you expect aluminum oxide to be a solid, liquid, or gas at room temperature? (b) Write a balanced chemical equation for the reaction of aluminum oxide with nitric acid

40. Practice Exercise (page 279) • Write the balanced chemical equation for the reaction of copper (II) oxide and sulfuric acid.

41. Nonmetals • Vary greatly in appearance • Have electron affinities that cause them to gain electrons when bonding with metals • Most nonmetal oxides are acidic • Ex: CO2(g) + H2O(l)  H2CO3(aq) P4O10 (s) + 6H2O (l)  4H3PO4 (aq)

42. Sample Exercise 7.9 • Write the balanced chemical equations for the reactions of solid selenium dioxide with (a) water (b) aqueous sodium hydroxide

43. Practice Exercise (page 281) • Write the balanced chemical equation for the reaction of solid tetraphosphorushexoxide with water.

44. Metalloids • Have properties between those of metals and nonmetals • Several (Si for example) are semiconductors and are used in computer chips

45. Group Trends: Alkali Metals • Soft, metallic solids • Silvery, metallic luster • High thermal and electrical conductivity • Low density and melting points • Very reactive (low first ionization energy) • Exist in nature only in compounds • Bonds with hydrogen to form hydrides • Can bond with oxygen to form metal peroxides (except lithium) • Some form superoxides (O2-) • Most cases form metal oxides • Each has characteristic flame color

46. Sample Exercise 7.10 • Write a balanced equation that predicts the reaction of cesium metal with (a) Cl2(g) (b) H2O (l) (c) H2(g)

47. Practice Exercise (page 283) • Write a balanced equation for the reaction between potassium metal and elemental sulfur.

48. Alkaline Earth Metals • Solids at room temperature • Harder, more dense, and higher melting points than alkali metals • Less reactive than alkali metals • Be & Mg are least reactive • Heavier alkaline earth metals give off a characteristic color when heated in a flame

49. Hydrogen • Nonmetal • Colorless, diatomic gas • Can be metallic under extreme pressure • High ionization energy • Less tendency to lose electrons than alkali metals • Share electrons when bonding with nometals to form molecular compounds • Reactions can be very exothermic • Reacts with active metals to from metal hydrides • Can lose electrons to form cation (H+)

50. Group 6A: Oxygen group • Nonmetallic  metallic • Oxygen is colorless gas • All other elements are solids • Allotropes (O2) and (O3) • Different forms of the same element in the same state • Sulfur…most common allotrope is S8 • Most sulfur in nature exists as metal sufide