Chapter 7: Periodic Properties of the Elements

1. Chapter 7: Periodic Properties of the Elements Alex Albert, Cleo Durham, Summer Ford, Kristen Johnson

2. Section 7.1: Development of the Periodic Table

3. Section 7.1, continued • Some elements appear in nature in their elemental form and were discovered thousands of years ago • Ex.: gold • Other elements are radioactive and intrinsically unstable, but recent technology has allowed us to discover them • Ex.: technetium

4. Section 7.1, continued • Majority of elements are not found in nature in their elemental form • Known elements: 31 in 1800 to 63 by 1865 • Mendeleev and Meyer discovered similar chemical and physical properties recur periodically when the elements are arranged in order of increasing atomic weight • Forerunners of modern periodic table

5. Section 7.1, continued • Mendeleev predicted existence and properies of Gallium (Ga) and Germanium (Ge), calling them “eka-aluminum” and “eka-silicon” • “Eka” = under • When discovered, properties closely matched predictions

6. Section 7.1, continued • Mosely developed concept of atomic numbers • Determined frequencies of x-rays emitted by elements as they were bombarded with high-energy electrons • Each element produces unique frequency and frequency increased with atomic mass • Arranged frequencies in order by assigned a whole number (atomic number)

7. Section 7.1, continued • This theory corrected problems created by organization according to atomic weight and led to discovery of other elements

8. Section 7.2: Effective nuclear charge

9. Section 7.2, continued • Coulomb’s Law: the strength of the interaction between two electrical charges depends on the magnitudes of the charges and on the distance between them • Applies to electron and nucleus • Force of attraction increases as nuclear charge increases and decreases as the electron moves father from the nucleus

10. Section 7.2, continued • Effective nuclear charge: the net positive charge experienced by an electron in a many-electron atom; this charge is not the nuclear charge because the is some shielding of the nucleus by the other electrons in the atom • Zeff < Z • Zeff= Z – S • Relationship between Zeff and number of protons in the nucleus, Z)

11. Section 7.2, continued • Inner electrons partially cancel the attraction between outer electrons and the nucleus • S – positive number called creencing constant; represents portion of nuclear charge screened from valence electrons • Value of S is usually close to number of core electrons in an atom • Screening of valence electrons in same shell slightly

12. Section 7.2, continued • Example: Na atom  condensed electron configuration: [Ne]3s¹ • Nuclear charge: 11+(2), inner core: 10 electrons (S) 11-10=1+ • Energies of orbitals with same “n” value increase with increasing “l” value • Carbon: 1s²2s²2p²  energy of 2p (l=1_ is slightly higher than 2s (l=0) • Radial probability fucntions: (ns<np<nd)

13. Section 7.2, continued • Effective nuclear charge increases moving across and row (period) of the table • Nuclear charge increases • Going down a column, effective nuclear charge experienced by valence electrons changes for less than it does across a row, increasing slightly

14. Section 7.3: Sizes of Atoms and ions

15. Section 7.3, continued • Apparent radii: the closest distances separating the nuclei of two atoms of the same element during collisions that occur while the element is in the gas phase • Radius is also called nonbonding atomic radius

16. Sections 7.3, continued • Bonding atomic radius: when two atoms of the same element are chemically bonded, chemists find the distance between the two nuclei and divide it in half • Bonding atomic radius: ½d, when d equals the distance between nuclei of bonded atoms • Generally accepted as the size of an atom • Also called covalent radius

17. Section 7.3, continued • Knowing atomic radii allows us to estimate the bond lengths between different elements in molecule • Ex.: Knowing the atomic radii of Cl and C allows chemists to find distance between nuclei of Cl and C when bonded

18. Section 7.3, continued • In regards to the trend of the size of atomic radii on the Periodic Table: • Within each column (group), atomic radius tends to increase form top to bottom. This trend results primarily from the increase in the principal quantum number (n) of the outer electrons. As we go down a column, the outer electrons have a greater probability of being father form the nucleus, causing the atom to increase in size. • Within each row (period), atomic radius tends to decrease form left to right. The major factor influencing this trend is the increase in the effective nuclear charge (Zeff) as we move across a row. The increasing effective nuclear charge steadily draws the valence electrons closer to the nucleus, causing the atomic radius to decrease.

19. Section 7.3, continued

20. Section 7.3, continued • Ionic radii can be determined from interatomic distances in ionic compounds • Cations are smaller than their parent atoms • The opposite is true for anions • When an electron is added to an atom to form an anion, the increased electron-electron repulsions cause the electrons to spread out more in space. • For ions carrying the same charge, size increases as we move down a column in the periodic table.

21. Section 7.3, continued • An isoelectronic series is a group of ions all containing the same number of electrons. • In any isoelectronic series we can list the members in order of increasing atomic number; therefore, nuclear charge increases as we move through the series • Because the number of electrons remains constant, the radius of the ion decreases with increasing nuclear charge, as the electrons are more strongly attracted to the nucleus.

22. Section 7.3, continued • For example: O−2 F− Na+ Mg+2 Al+3 10 electrons each Increasing nuclear charge  Decreasing atomic radius  • Note: Oxygen, the largest ion in this series, has the lowest atomic number (8), while Aluminum the smallest ion, has the highest atomic number (13).

23. Section 7.4: ionization energy

24. Section 7.4, continued • Ionization energy of an atom or ion is the minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion. • The first ionization energy, I1, is the energy needed to remove the first electron from a neutral atom. • Ex.: Na(g)  Na+ + e- • The second ionization energy, I2, is the energy needed to remove the second electron. • Ex.: Na+(g)  Na2+ + e- • The greater the ionization energy, the more difficult it is to remove an electron.

25. Section 7.4, continued • Periodic Trends in the First Ionization Energies • Within each row (period) on the table, I1 generally increases with increasing atomic number. The alkali metals show the lowest ionization energy in each row, and the noble gases show the highest. • Within each column (group) of the table, the ionization energy generally decreases with increasing atomic number. For example, the ionization energies of the Noble Gases are: He > Ne > Ar > Kr > Xe • The s- and p-block elements show a larger range of values for I1 than doe the transition-metal elements. Generally, the ionization energies of the transition metals increase slowly as we proceed from left to right in a period.

26. Section 7.4, continued

27. Section 7.4, continued • In general, smaller atoms have higher ionization energies.

28. Section 7.5: Election Affinities

29. Section 7.5, continued • Most atoms can gain e⁻ to form negatively charged ions.   • Electron Affinity- the energy change that occurs when an e⁻ is added to a gaseous atom (because it measures the attraction, or affinity, of the atom for the added e⁻) • Most atoms release energy when an e⁻ is added

30. Section 7.5, continued • Ex. Electron Affinity: Cl(g) + e⁻ → Cl⁻(g) ∆E= -349kJ/mol

31. Section 7.5, continued • Ionization energy measures the ease with which an atom loses an e⁻. • Electron Affinity measures the ease with which an atom gains an e⁻. • The greater the attraction b/t a given atom and an added e⁻, the more negative the atoms electron affinity will be.

32. Section 7.6: Metals, Nonmetals and metalloids

33. Section 7.6, continued • Broad groupings of the elements • Metallic Character: how much an element exhibits the physical and chemical properties characteristic of metals (conductivity, luster, malleability, ductility, etc)

34. Section 7.6, continued • Metals • Shiny luster, conductors, generally malleable and solid • Tend to have low ionization energies and form positive ions pretty easily • 1st ionization energy is an indicator of metal or nonmetal • Compounds of metals with nonmetals tend to be ionic substances

35. Section 7.6, continued • Nonmetals • Very in appearance (generally poor conductors, not lustrous) • Tend to gain electrons when they react with metals b/c their electron affinities • Compounds composed entirely of nonmetals are typically molecular substances

36. Section 7.6, continued • Metalloids • Properties b/t those of metals and nonmetals • Ex.: Silicon looks like metal but is brittle rather than malleable and is a poor conductor of heat and electricity • Compounds can have characteristics of the compounds of metals of nonmetals (all depends upon the specific compound)

37. Section 7.7: Group Trends for the Active Metals

38. Group 1A: The Alkali Metals

39. Section 7.7, continued • Alkali metals • Soft metallic solids • Silvery luster • High thermal and electrical conductivities • Low densities and melting points (D=g/cm3) • Very reactive

40. Section 7.7, continued • Alkali metals • Tarnish rapidly on exposure to air • One electron in their valence electron shell • From ions of 1+ charge • In each row of the periodic table, the alkali metals have the lowest I value. It reflects how easy outer electron can be removed • Alkali metals lose electrons to form compounds with nonmetals

41. Section 7.7, continued

42. Section 7.7, continued • Reactions of alkali metals and water produce H2 and an alkali metal hydroxide solution • 2M (s) + 2H2O (l)  2MOH (aq) + H2(g) • Reactions with Li and water is slow and steady • Reactions with Na and water is vigorous, creating enough heat to melt the sodium • Reactions with K and water is violent and the heat produced is enough to ignite the hydrogen evolved • Reactions of soluble metal oxides and water such as Li2O, form hydroxide ions

43. Group 7A: The Halogens

44. Section 7.7, continued • Halogens • Nonmetals • Very reactive • Non-polar • Require one electron to complete their shell • Slightly soluble, dissolve in water to form hydrohalide acids

45. Section 7.7, continued • Halogens • Melting/boiling points increase with increasing atomic # • React with most metals to form ionic halides; react with hydrogen to form gaseous hydrogen halide compounds • Negative affinities; tendencies to gain electrons from other elements to form halide X- where X= halogen element • Diatomic molecules (F2, Cl2, Br2, I2)

46. Section 7.7, continued

47. Section 7.7, continued • Trends in the oxides of the Period 3 elements • At the left hand side the elements have relatively low ionization energies so they bond to form ionic compounds (Na and Mg). • They lose their valence electron to be cations • The oxides of the elements are ionic and have an oxide ion • Then the oxide ion can form a bond to hydrogen ions and as a result these ionic oxides act as bases dissolving in water • When ionization energy becomes too great form cation formation the elements make covalent bonds (C and Si) • At the far right of the period (except for the noble gases) the elements can and do hold covalent bonds by sharing electrons with non-metals (S and Cl). These elements also have the option of forming ionic compounds with metals in which they will exist as anions. (S2- and Cl-)

48. Section 7.7, continued • Trends in the oxides of Period 3 elements • When ionization energy becomes too great form cation formation the elements make covalent bonds (C and Si) • At the far right of the period (except for the noble gases) the elements can and do hold covalent bonds by sharing electrons with non-metals (S and Cl). These elements also have the option of forming ionic compounds with metals in which they will exist as anions. (S2- and Cl-)

49. Section 7.7, continued

50. PRACTICE QUESTIONS