The Periodic Table

1. The Periodic Table Metals, Nonmetals, & Metalloids, Families, Periodic Trends

2. Periodic Properties of the Elements • Modern periodic table arranged in order of increasing atomic number. • 1869: Dmitri Mendeleev and Lothar Meyer publish identical tables. Mendeleev gets the credit and becomes “Father of Periodic Table”. • 1913: Henry Moseley develops concept of atomic numbers.

4. Characteristics of Metals • Most elements on the periodic table are metals • Have a shiny luster; most are silvery • All are solids at room temperature except Hg • Are malleable (thin sheets) and ductile (wires) • Many have a very high melting point (1900°C for chromium) • Good conductors of heat and electricity • Have low electron affinity; form cations • Compounds of metals with nonmetals are ionic

5. Metal Oxides Most metal oxides are basic because oxide ion reacts with water to form hydroxide ion. O2-(aq) + H2O(l) 2OH-(aq) Metal oxides that dissolve in water form metal hydroxides: CaO(s) + H2O(l) Ca(OH)2 (aq) Metal oxides that react with acids form salts and water: NiO(s) + 2HCl(aq) NiCl2(aq) + H2O(l)

6. Characteristics of Nonmetals • Do not have luster; various colors • Poor conductors of heat and electricity • Melting points are generally lower than those of metals • May be solids, liquids, or gases at room temp • Have high electron affinity; form anions • Compounds of nonmetals with metals are ionic • Compounds with only nonmetals are molecular

7. Nonmetal Oxides Most nonmetal oxides are acidic. Nonmetal oxides that dissolve react with water to form acids: CO2 (g) + H2O(l) H2CO3 (aq) Like acids, most nonmetal oxides dissolve in basic solutions to form salts and water. CO2 (g) + 2NaOH(aq) Na2CO3 (aq) + H2O(l)

8. Characteristics of Metalloids • Include the elements B, Si, Ge, As, Sb, Te, At • Form the division line between metals and nonmetals on the periodic table • Have properties intermediate between metals and nonmetals • Metalloids such as Si and Ge are electrical semiconductors and are used in integrated circuits and computer chips.

9. Group 1A: The Alkali Metals • Exist at room temperature as soft, silver solids • Na and K are abundant in seawater, Earth, and biological systems • Have low densities and melting points • Form ions with a 1+ charge • Very reactive; combine with most nonmetals • React vigorously with water in an exothermic reaction • Emit characteristic colors when placed in a flame

10. Group 2A: Alkaline Earth Metals • Harder and more dense than Group 1A • Have higher melting points than Group 1A • Form cations with a 2+ charge • Mg and Ca are essential for living organisms • Emit characteristic colors when placed in a flame

11. Group 6A: Chalcogens • Group has elements with nonmetallic and metallic character • Oxygen has two allotropes: O2 and O3 (ozone) • Oxygen readily oxidizes other elements • Sulfur’s allotrope (S8) is a yellow solid composed of eight-membered rings • Oxygen and sulfur form anions with a 2- charge

12. Group 7A: Halogens • Greek word meaning “salt formers” • Main halogens are F, Cl, Br, I • Are diatomic at room temperature: F2 and Cl2 are gases, Br2 is liquid, I2 is solid • Form halide ions with a 1- charge • Fluorine is most reactive

13. Group 8A: Noble Gases • All are monoatomic gases at room temperature • Have stable electron configurations and are exceptionally unreactive • Some compounds of Xe and F were synthesized in the 1960’s

14. Hydrogen • Does not truly belong to any group • Is a nonmetallic diatomic gas at room temp • Can be metallic at extreme pressures • Usually forms covalent bonds • Can form hydrogen ion (1+ charge) or hydride ion (1- charge)

15. Coulomb’s Law of Attraction So…the force of attraction between the nucleus and the electron depends on the net nuclear charge acting of the electronand the average distance between the nucleus and the electron. F = kQ1Q2 d2 • The strength of interaction between two electrical charges depends on the magnitude of the charges and the distance between the two.

16. Effective Nuclear Charge (Zeff) • In an atom, each electron is simultaneously attracted to the nucleus (+ charge) and repelled by other electrons (- charge). • We can estimate the energy of an individual electron in the average electric field created by the nucleus and the electron density of the other electrons. • We look at the effective nuclear charge (Zeff) located at the nucleus.

17. How Do We Determine Effective Nuclear Charge (Zeff)? Zeff = Z - S • Zeff = charge acting on electron by nucleus • Z = number of protons in nucleus • S = Average number of electrons between nucleus and specific electron. S can be a non-integer.

18. What is Electron Shielding? • Electron density of core e-’s shield or screen outer e-’s from the full charge of the nucleus. • Electrons in the same shell do little shielding but do repulse each other. • In a multielectron system an e- in any orbital will partially shield an e- in any other orbital.

19. What Determines Effective Nuclear Charge (Zeff)? • The Zeff experienced by outer electrons is determined primarily by the difference between the charge on the nucleus and the charge of the core electrons. • The force of attraction between the electron and the nucleus increases as the nuclear charge increases and decreases as the electron moves farther from the nucleus.

20. Estimating Zeff We can roughly estimate Zeff for an outer electron in a magnesium atom. Zeff = Z - S Z = number of protons = 12+ S = avg. # of core electrons = 10 Zeff = 12 - 10 = 2+ charge

22. How Rough is our Estimate of Magnesium’s Zeff? • The actual effective nuclear charge for an outer magnesium electron is about 3.3+ compared with our calculated 2+ charge. • Our equation underestimates Zeff because it ignores the times when the outer electron may be inside the core, closer to the nucleus. This penetration causes the Zeff to be higher.It also ignores the electron-electron repulsions.

23. Effective nuclear charge steadily increases as we go from left to right across a period. WHY? The charge increases because the number of protons increases but the number of core electrons (shielding) stays the same. Effective nuclear charge increases slightly as we go down a family. WHY? Larger cores are less able to shield the outer electrons. Periodic Table Trend 1: Effective Nuclear Charge Remember: Electron shielding is constant across a period because the core electrons remain constant! Note: This trend down a family is much less important that the trend across a period.

24. Periodic Trend 2: Atomic Radius • Atoms do not have sharply defined boundaries (think electron cloud) but an atomic radius is known for most elements. • Nonbonding atomic radius (van der Waals radius): the radius of an atom as defined by the closest distance separating its nucleus from the nucleus of another atom during a collision. 0.5 d • Bonding atomic radius (covalent radius): theradius of an atom as defined by the distance separating its nucleus from the nucleus of another atom to which it is chemically bonded.

25. Atomic radius decreases up a family (bottom to top of periodic table). WHY? Radius decreases with decreasing n because the electrons spend more time closer to the nucleus, thereby decreasing atomic size. Atomic radius decreases across period from left to right. WHY? Shielding stays constant across a period so Zeff increases. Greater Zeff draws the electrons in closer to the nucleus, thereby decreasing atomic size. Periodic Trend 2: Atomic Radius

27. CATIONS Radius of a cation is smaller than its parent atom. WHY? Outermost electrons leave an atom to form a cation. b. New outermost electrons are in a lower energy level, closer to the nucleus. ANIONS Radius of an anion is larger than its parent atom. WHY? a. Electrons are added to outer shell to form an anion. b. More electrons means more electron-electron repulsion. Periodic Trend 3: Ionic Radius

29. Cation radii decrease up a family (bottom to top). WHY? Atomic radii decrease up a family, and cations are smaller than their parent atoms. Larger atoms at the bottom of the table have larger ions than smaller atoms at the top of the table. Anion radii decrease up a family (bottom to top). WHY? Atomic radii decrease up a family, and anions are larger than their parent atoms. Larger atoms at the bottom of the table have larger ions than smaller atoms at the top of the table. Periodic Trend 3: Ionic Radius Trend Up a Family

31. Periodic Trend 3: Ionic Radius Trend Across a Period • The ionic radius trend changes across a period from left to right as we move from metals to nonmetals. • The radii of metallic ions decrease while the radii of nonmetallic ions increase across a period. • This variation between cations and anions causes a wave-like pattern on the table.

32. Isoelectronic Series of Ions • Isoelectronic ions have the same number of electrons. Example: O2-, F-, Na+, Mg2+, Al3+ all have 10 electrons. Increasing nuclear charge  O2-, F-, Na+, Mg2+, Al3+ Decreasing ionic radius  • In an isoelectronic series of ions, the ion with the most negative charge has the largest radius.

33. Ionization Energy • Ionization Energy is the minimum amount of energy needed to remove an electron from the ground state of an isolated gaseous atom or ion. First Ionization Energy (I1) removes the first electron from a neutral atom: Na(g)  Na+(g) + 1e- Second Ionization Energy (I2) removes the second electron: Na+(g)  Na2+(g) + 1e- • Formation of cations The greater the ionization energy, the more difficult it is to remove an electron.

34. Variations in Successive Ionization Energies • Ionization energy for removal of successive electrons: I1 < I2 < I3 • There is a sharp increase in • ionization energy as inner shell • electrons are removed. • Every element shows a large • increase in ionization energy • when electrons are removed • from its noble gas core.

36. WHY? As atomic radius decreases, it is harder to remove electrons because they are more attracted to the nucleus. WHY? As Zeff increases and the atom gets smaller, it is harder to remove electrons because they are more attracted to the nucleus. Periodic Trend 4: Ionization Energy Ionization energy generally increases as we go up a family (bottom to top). Ionization energy generally increases as we go from left to right across a period.

39. What Is Electron Affinity? • Ionization energy measures how easily an atom loses an electron and electron affinity measures how easily an atom gains an electron. • Electron affinity is the energy change that occurs when an electron is added to a gaseous atom. It measures the attraction of the atom for the added electron.

40. Basic Electron Affinity Concepts • For most atoms, energy is released when an electron is added. Therefore, electron affinities are usually negative values. KABOOM! Example: Cl(g) + 1e- Cl-(g) ∆E = -349 kJ/mol • The greater the attraction between an atom and an added electron, the more negative the atom’s electron affinity will be.

41. What Does It Mean If The Electron Affinity Is Positive? • Some elements such as noble gases have positive electron affinities. Example: Ar(g) + 1e- Ar- (g) ∆E > 0 kJ/mol • This means that the argon anion is higher in energy than the separated atom and electron. Because ∆E > 0, the argon ion is unstable and does not form. Ar-

42. Electron affinity does not change greatly as we move up a family(bottom to top). WHY? Increasing electron-nucleus attraction is counterbalanced by higher electron-electron repulsions. Electron affinity generally becomes increasingly negative as we go from left to right across a period. Halogens have the most negative electron affinites. Noble gases all have positive electron affinities. WHY? Halogens are one electron shy of a very stable noble gas configuration. Periodic Trend 5: Electron Affinity

44. Extra Questions From Chapter • 1. How does a molecule of oxygen differ from a molecule of sulfur in its most common form at room temp.? • 2. What are the names and formulas of the allotropes of oxygen? • 3. What are the names and formulas of the allotropes of carbon? • 4. Do you expect astatine to a solid, liquid, or gas at room temp.? • 5. Which is a better conductor of electricity: tellurium or iodine? • 6. The species Na, Mg+, Al2+, and Si3+ are isoelectronic. For which of these will the Zeff be the smallest?

45. More Questions… • 7. Arrange in order of increasing atomic radius: Cs, K, Rb • 8. Select the two ions that are isoelectronic with each other: K+, Rb+, Ca2+, Se2-, Sc3+ • 9. Write the equation for the 2nd ionization of aluminum. • 10. Write the e- configuration for Sc3+. • 11. Write the e- configuration for Mn3+. • 12. As metallic character increases, first ionization energy ___________________. • 13. List three characteristics of metals. • 14. Which oxide is the most acidic: magnesium oxide, aluminum oxide, or sulfur trioxide?

46. Still More Questions… • 15. What is the effective nuclear charge for an outer aluminum electron? • 16. Who is the “Father of the Periodic Table”? • 17. The periodic table is arranged in order of increasing _______________. • 18. T or F: In any particular compound the actual radius of an atom may deviate from the average calculated radius. • 19. Who developed the concept of atomic numbers? • 20. T or F: A more negative value for electron affinity denotes a greater electron affinity.