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Chapter 6 The Periodic Table

Chapter 6 The Periodic Table. Why are the elements cobalt phosphorus and sulfur essential for maintaining law and order?. Because when they are put together, they’re Co P S. 6.1 Organizing the Elements. OBJECTIVES: Explain how elements are organized in a periodic table

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Chapter 6 The Periodic Table

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  1. Chapter 6The Periodic Table Why are the elements cobaltphosphorusandsulfuressential for maintaining law and order? Because when they are put together, they’re CoPS

  2. 6.1 Organizing the Elements • OBJECTIVES: • Explain how elements are organized in a periodic table • Compare early and modern periodic tables • Identify three broad classes of elements

  3. Section 6.1Organizing the Elements • A few elements, such as gold and copper, known for thousands of years • Yet only about 13 had been identified by the year 1700. • As more elements discovered, chemists realized a way to organize the elements was needed.

  4. 6.1 Organizing the Elements • Chemists used properties of elements to sort them into groups. • In 1829 J. W. Dobereiner arranged elements into triads – groups of three elements with similar properties • One element in each triad had properties intermediate of other two • Worked well for many elements, but not all

  5. Mendeleev’s Periodic Table • By mid-1800s, about 70 elements known • Dmitri Mendeleev, a Russian chemist and teacher, arranged elements in order of increasing atomic mass • Mendeleev noticed that if he grouped elements a certain way, certain chemical characteristics appeared periodically on the table • Thus, the Periodic Table

  6. Mendeleev • Mendeleev left “blanks” on his table where he thought undiscovered elements should fit • When some of these elements were discovered, many of his predictions proved to be spot on • But, there were problems: • Co and Ni; Ar and K; Te and I • He’d switched them on his table because he didn’t know about isotopes

  7. A better arrangement • Mendeleev’s mistake: arranging elements based on atomic mass instead of z • so he’d switched Co and Ni because the atomic mass of Co is slightly higher than Ni, due to Co having heavier isotopes (more neutrons) • In 1913, Henry Moseley, a British physicist, arranged elements according to increasing atomic number • The symbol, atomic number & mass are basic items included on most tables

  8. Spiral Periodic Table

  9. The Periodic Law says: • When elements are arranged in order of increasing atomic number, a periodic repetitionof physical and chemical properties occurs • Horizontal rows = periods • There are 7 periods • Vertical columns = groups (or families) • Similar physical & chemical props • Especially main block elements(groups 1A - 8A)

  10. Classes of Elements • Metals: electrical conductors, have luster, ductile, malleable • Nonmetals: brittle non-lustrous solids, or gases, one liquid (Br), all poor conductors of heat & electricity • Metalloids: border zig-zag line on both sides and have properties intermediate between metals and nonmetals

  11. Checking Understanding • What did Mendeleev base his table of elements upon? • What caused Mendeleev to mistakenly switch Co and Ni; Ar and K; Te and I on his original table? • Complete the following table (Note charges on ions):

  12. 6.2 Classifying the Elements • OBJECTIVES: • Describe the information provided in a periodic table • Classify elements based on electron configurations • Distinguish representative elements (main block) and transition metals

  13. Squares in a Periodic Table • The periodic table displays symbols and names of elements, along with information about the structure of their atoms: • Atomic number and atomic mass • Symbols usually color coded • solid; gas; liquid

  14. Basedontheaverageatomicmassand theatomicnumber, whatislikelythemostcommonisotope of Na?

  15. Groups of Elements - Families • Group 1A – alkali metals • Forms a “base” (or alkali) when reacting (very violently!) with water • Na + 2 H2O  Na+ + 2 OH− + H2 • Group 2A – alkaline earth metals • Also form bases with water • do not dissolve well and melt at very high T, hence usually solid (“earth”) • Group 7A – halo-gens (= “salt-forming”)

  16. Electron Configurations in Groups • Elements can be sorted into 4 different groupings based on their electron configurations: • Noble gases • Representative elements • Transition metals • Inner transition metals Let’s now take a closer look at these.

  17. Electron Configurations in Groups • Main Block Elements = Groups 1A - 8A • Representative ElementsGroups 1A - 7A • Display wide range of properties, thus a good “representative” sample of all matter • There are metals, nonmetals, and metalloids; Many are solid, others are gases or liquids • Their outermost electrons are in s and psubshells that are NOT filled • Noble Gases: so called because they rarely take part in a reaction; very stable (exist as free atoms in nature) • Noble gases have their outer s and p sublevels completely full (the most stable arrangement)

  18. Electron Configurations in Groups • Transition (d-Block) metals in the “B” columns of the U.S. style periodic table • Electron configuration has the outer s sublevel full (usually), and all or part of d subshell filled • Block “transitions” between the metal and nonmetal area of table • Examples include familiar metals: gold, copper, silver, iron

  19. Electron Configurations in Groups • Inner Transition Metals located below main body of the table, in two rows • Electron configuration has the outer s sublevel full, and f sublevel is partly to completely filled • Formerly called “rare-earth” elements, but this is not true because some are quite abundant • We won’t be concerned with these much in this class

  20. 8A 1A • Main Block in color • Groups 1A-7A are representative elements 2A 3A 4A 5A 6A 7A

  21. These are called the inner transition elements, and they belong here The group B are called the transition elements

  22. Group 1A are the alkali metals (but NOT H) Group 2A are the alkaline earth metals H

  23. Group 8A are the noble gases • Group 7A is called the halogens

  24. H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87 Do you notice any similarity in the configurations of the alkali metals? 1s1 1s22s1 1s22s22p63s1 1s22s22p63s23p64s1 1s22s22p63s23p64s23d104p65s1 1s22s22p63s23p64s23d104p65s24d10 5p66s1 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1

  25. He Do you notice any similarity in the configurations of the noble gases? 2 1s2 1s22s22p6 1s22s22p63s23p6 1s22s22p63s23p64s23d104p6 1s22s22p63s23p64s23d104p65s24d105p6 1s22s22p63s23p64s23d104p65s24d10 5p66s24f145d106p6 Ne 10 Ar 18 Kr 36 Xe 54 Rn 86

  26. Elements in the s - blocks s1 • Alkali metals all end in s1 • Alkaline earth metals all end in s2 • really should include He, but it fits better in a different spot, since He has the properties of the noble gases, and has a full outer level of electrons. s2 He

  27. Transition Metals - d block Note exceptions to electron configuration rules s1 d5 s1 d10 d1 d2 d3 d5 d6 d7 d8 d10

  28. The P-block p1p2 p3 p4 p5p6

  29. 1 2 3 4 5 6 7 • Each row (or period) is the energy level for s and p orbitals (d block is one level lower in each row). Period Number 3d  4d 

  30. 1 2 3 4 5 6 7 • f orbitals start filling at 4f, and are 2 less than the period number 4f 5f

  31. 6.3 Periodic Trends • OBJECTIVES: • Describe and explaintrends among the elements for atomic size • Explain how ions form • Describe and explain periodic trends for first ionization energy, ionic size • Describe the periodic trend for electronegativity

  32. Trends in Atomic Size • How do we measure? • electron clouds of atoms don’t have a definite edge • Chemists get around this by measuring more than 1 atom at a time

  33. Atomic Size } • Measure the Atomic Radius - this is half the distance between the two nuclei of a diatomic molecule. Radius

  34. 3 Factors Affect Periodic Trends 1. Energy Level • Electrons in higher energy levels are farther away from the nucleus • Each successive level builds on top of lower ones 2. Nuclear Charge (# protons) • More + charge pulls electrons in closer. (+ and – attract each other) 3. Shielding effect • “core” electrons near nucleus shield outer electrons from protons’ “pull” • The higher the energy level  greater shielding

  35. #1. Atomic Size - Group trends H • As we go down a group, each atom has another energy level (= more electrons, farther from nucleus) • so the atoms get bigger. Li Na K Rb

  36. #1. Atomic Size - Period Trends • Going from left to right across a period, the size getssmaller. • Electrons are in the same energy level. • But, there is more nuclear charge. • Outermost electrons are pulled closer. Na Mg Al Si P S Cl Ar

  37. Rb K Period 2 Na Li Atomic Radius (pm) Kr Ar Ne H Atomic Number 3 10

  38. Ions • Some compounds are composed of particles called “ions” • An ion is a atom (or group of atoms) that has a positive or negative charge • Atoms are neutral because the number of protons equals electrons • Positive and negative ions are formed when electrons are transferred (lost or gained) between atoms

  39. Ions • Metals tend to LOSE electrons, from their outer energy level • Sodium loses one: there are now more protons (11) than electrons (10), and thus a positively charged particle is formed = “cation” • Charge is written as a number followed by a + or – sign: Na1+ orNa+ • Now named a “sodium ion”

  40. Ions • Nonmetals tend to GAIN one or more electrons • Oxygen often gains two electrons • Protons (8) no longer equals electrons (10), so a charge of -2 • O2– is re-named a “oxide ion” • Negative ions are called “anions”

  41. #2. Trends in Ionization Energy • Ionization energy is the amount of energy required to completely remove an electron (from a gaseous atom) • Removing one electron makes a 1+ ion • The energy required to remove only the first electron is called the first ionization energy • Energy needed to remove a second electron = second ionization energy • Remove a third = third ionization energy • And so on…

  42. Ionization Energy • The second ionization energy is always greater than first IE • There are more protons than electrons after the first is removed, so the nucleus pulls the rest closer • The third IE (energy to remove a third electron) is even greater • Removing an electron from a 2+ ion

  43. Table 6.1, p. 173 Symbol First Second Third 5247 7297 1757 2430 2352 2857 3391 3375 3963 1312 2731 520 900 800 1086 1402 1314 1681 2080 HHeLiBeBCNO F Ne 11810 14840 3569 4619 4577 5301 6045 6276

  44. Symbol First Second Third 11810 14840 3569 4619 4577 5301 6045 6276 5247 7297 1757 2430 2352 2857 3391 3375 3963 1312 2731 520 900 800 1086 1402 1314 1681 2080 HHeLiBeBCNO F Ne Why did these values increase so much?

  45. What factors determine IE • Greater nuclear charge = greater IE • More protons, more difficult to remove e- • Greater distance from nucleus decreases IE • Farther from “pull” of protons + shielding effect of e- in lower energy levels • Filled and half-filled orbitals have lower energy, so achieving them is easier, lower IE.

  46. Shielding • Nucleus has weak attraction to electron in outermost energy level due to repulsive forces of all the other energy levels in between. • Removing a second e– from the same period requires only a little more energy because shielding is the same

  47. Symbol First Second Third 6912 7732 2744 39634565 14501816 2080 496 738 578 NeNaMgAl Why did these values increase so much? The big jumps in IE seen above occur because that second electron removed from Na and the third removed from Mg are being taken from a lower energy level (a core electron rather than a valence electron)

  48. Ionization Energy - Group trends • As you go down a group, the first IE decreases because... • The electron is further away from the attraction of the nucleus, and • There is more shielding.

  49. Ionization Energy - Period trends • All the atoms in the same period have the same energy level • = same shielding • BUT, increasing nuclear charge • So IE generally increases from left to right • Exceptions occur (of course), but can be explained in light of electron configurations

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