The periodic table is a tabular arrangement of chemical elements that are arranged by increasing atomic number and groups elements according to recurring properties. The rows are arranged so that metals are on the left side of the table and nonmetals are on the right side. The columns are called groups.
1. “The Periodic Table”
2. Organizing the Elements OBJECTIVES: • Explain how elements are organized in a periodic table.
3. Organizing the Elements OBJECTIVES: • Compare early and modern periodic tables.
4. Organizing the Elements OBJECTIVES: • Identify three broad classes of elements.
5. Organizing the Elements A few elements, such as gold and copper, have been known for thousands of years - since ancient times Yet, only about 13 had been identified by the year 1700. As more were discovered, chemists realized they needed a way to organize the elements.
6. Organizing the Elements Chemists used the 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 the other two elements
7. Mendeleev’s Periodic Table By the mid-1800s, about 70 elements were known to exist Dmitri Mendeleev – a Russian chemist and teacher Arranged elements in order of increasing atomic mass Thus, the first “Periodic Table”
8. Mendeleev He left blanks for yet undiscovered elements • When they were discovered, he had made good predictions But, there were problems: • Such as Co and Ni; Ar and K; Te and I
9. A better arrangement In 1913, Henry Moseley – British physicist, arranged elements according to increasing atomic number The arrangement used today The symbol, atomic number & mass are basic items included- textbook page 162 and 163
10.
11. Another possibility: Spiral Periodic Table
12. The Periodic Law says: When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties. Horizontal rows = periods • There are 7 periods Vertical column = group (or family) • Similar physical & chemical prop. • Identified by number & letter (IA, IIA)
13. Areas of the periodic table Three classes of elements are: 1) metals, 2) nonmetals, and 3) metalloids 1) Metals: electrical conductors, have luster, ductile, malleable 2) Nonmetals: generally brittle and non-lustrous, poor conductors of heat and electricity
14. Areas of the periodic table Some nonmetals are gases (O, N, Cl); some are brittle solids (S); one is a fuming dark red liquid (Br) Notice the heavy, stair-step line? 3) Metalloids: border the line-2 sides • Properties are intermediate between metals and nonmetals
15. Classifying the Elements OBJECTIVES: • Describe the information in a periodic table.
16. Classifying the Elements OBJECTIVES: • Classify elements based on electron configuration.
17. Classifying the Elements OBJECTIVES: • Distinguish representative elements and transition metals.
18.
19. Groups of elements - family names Group IA – alkali metals • Forms a “base” (or alkali) when reacting with water (not just dissolved!) Group 2A – alkaline earth metals • Also form bases with water; do not dissolve well, hence “earth metals” Group 7A – halogens • Means “salt-forming”
20. Electron Configurations in Groups Elements can be sorted into 4 different groupings based on their electron configurations: 1) Noble gases Let’s 2) Representative elements now take a 3) Transition metals closer look at 4) Inner transition metals these.
21. Electron Configurations in Groups 1) Noble gases are the elements in Group 8A (also called Group18 or 0) • Previously called “inert gases” because they rarely take part in a reaction; very stable = don’t react • Noble gases have an electron configuration that has the outer s and p sublevels completely full
22. Electron Configurations in Groups 2) Representative Elements are in Groups 1A through 7A • Display wide range of properties, thus a good “representative” • Some are metals, or nonmetals, or metalloids; some are solid, others are gases or liquids • Their outer s and p electron configurations are NOT filled
23. Electron Configurations in Groups 3) Transition metals are in the “B” columns of the periodic table • Electron configuration has the outer s sublevel full, and is now filling the “d” sublevel • A “transition” between the metal area and the nonmetal area • Examples are gold, copper, silver
24. Electron Configurations in Groups 4) Inner Transition Metals are located below the main body of the table, in two horizontal rows • Electron configuration has the outer s sublevel full, and is now filling the “f” sublevel • Formerly called “rare-earth” elements, but this is not true because some are very abundant
25. Elements in the 1A-7A groups 1A 8A are called the representative 2A elements 3A 4A 5A 6A 7A outer s or p filling
26. The group B are called the transition elements These are called the inner transition elements, and they belong here
27. Group 1A are the alkali metals (but NOT H) Group 2A are the alkaline earth metals H
28. Group 8A are the noble gases Group 7A is called the halogens
29. H 1s1 1 Do you notice any similarity in these configurations of the alkali metals? Li 1s22s1 3 Na 1s22s22p63s1 11 K 1s22s22p63s23p64s1 19 Rb 1s22s22p63s23p64s23d104p65s1 37 Cs 1s22s22p63s23p64s23d104p65s24d10 55 5p66s1 87 1s22s22p63s23p64s23d104p65s24d105p66s 24f145d106p67s1
30. He Do you notice any similarity in the 1s 2 2 configurations of the noble gases? Ne 1s22s22p6 10 Ar 1s22s22p63s23p6 18 Kr 1s 2s 2p 3s 3p 4s 3d 4p 2 2 6 2 6 2 10 6 36 1s22s22p63s23p64s23d104p65s24d105p6 Xe 54 1s22s22p63s23p64s23d104p65s24d10 Rn 5p66s24f145d106p6 86
31. Elements in the s - blocks He 2 s 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.
32. Transition Metals - d block Note the change in configuration. s1 s1 d1 d2 d3 d 5 d5 d6 d7 d8 d10 d10
33. The P-block p1 p2 p6 p3 p4 p 5
34. F - block Called the “inner transition elements” f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
35. 1 2 3 4 5 6 7 Each row (or period) is the energy level for s and p orbitals.
36. The “d” orbitals fill up in levels 1 less than the period number, so the first d is 3d even though it’s in row 4. 1 2 3d 3 4 5 6 7
37. 7 4f 5f f orbitals start filling at 4f, and are 2 less than the period number
38. Periodic Trends OBJECTIVES: • Describe trends among the elements for atomic size.
39. Periodic Trends OBJECTIVES: • Explain how ions form.
40. Periodic Trends OBJECTIVES: • Describe periodic trends for first ionization energy, ionic size, and electronegativity.
41. Trends in Atomic Size First problem: Where do you start measuring from? The electron cloud doesn’t have a definite edge. They get around this by measuring more than 1 atom at a time.
42. Atomic Size } Radius Measurethe Atomic Radius - this is half the distance between the two nuclei of a diatomic
43. ALL Periodic Table Trends Influenced by three factors: 1. Energy Level • Higher energy levels are further away from the nucleus. 2. Charge on nucleus (# protons) • More charge pulls electrons in closer. (+ and – attract each other) 3. Shielding effect (blocking effect?)
44. What do they influence? Energy levels and Shielding have an effect on the GROUP ( ) Nuclear charge has an effect on a PERIOD ( )
45. #1. Atomic Size - Group trends As we increase H the atomic Li number (or go down a group). . . Na each atom has K another energy level, so the atoms get Rb bigger.
46. #1. Atomic Size - Period Trends Going from left to right across a period, the size gets smaller. 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
47. Rb K Period 2 Atomic Radius (pm) Na Li Kr Ar Ne H 3 10 Atomic Number
48. Ions Some compounds are composed of particles called “ions” • An ion is an 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
49. 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” • The charge is written as a number followed by a plus sign: Na1+ • Now named a “sodium ion”
50. Ions Nonmetals tend to GAIN one or more electrons • Chlorine will gain one electron • Protons (17) no longer equals the electrons (18), so a charge of -1 • Cl1- is re-named a “chloride ion” • Negative ions are called “anions”
51. #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.
52. Ionization Energy The second ionization energy is the energy required to remove the second electron. • Always greater than first IE. The third IE is the energy required to remove a third electron. • Greater than 1st or 2nd IE.
53. Table 6.1, p. 173 Symbol First Second Third H 1312 He 2731 5247 Li 520 7297 Be 900 1757 11810 B 800 2430 14840 C 1086 2352 3569 N 1402 2857 4619 O 1314 3391 4577 F 1681 3375 5301 Ne 2080 3963 6045 6276
54. Symbol First Second Third H 1312 Why did these values 2731 5247 increase so much? He Li 520 7297 11810 Be 900 1757 14840 B 800 2430 3569 C 1086 2352 4619 N 1402 2857 4577 O 1314 3391 5301 F 1681 3375 6045 Ne 2080 3963 6276
55. What factors determine IE The greater the nuclear charge, the greater IE. Greater distance from nucleus decreases IE Filled and half-filled orbitals have lower energy, so achieving them is easier, lower IE. Shielding effect
56. Shielding The electron on the outermost energy level has to look through all the other energy levels to see the nucleus. Second electron has same shielding, if it is in the same period
57. 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.
58. 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 at full and 1/2 full orbitals.
59. He He has a greater IE First Ionization energy than H. Both elements have H the same shielding since electrons are only in the first level But He has a greater nuclear charge Atomic number
60. First Ionization energy He Li has lower IE than H more shielding H further away These outweigh Li the greater nuclear charge Atomic number
61. First Ionization energy He Be has higher IE than Li same shielding H Be greater nuclear charge Li Atomic number
62. He B has lower IE First Ionization energy than Be same shielding H greater nuclear Be charge By removing an B Li electron we make s orbital half-filled Atomic number
63. First Ionization energy H He Li Be B C Atomic number
64. First Ionization energy He N H C Be B Li Atomic number
65. First Ionization energy He Oxygen breaks N the pattern, because H Be C O removing an electron leaves B it with a 1/2 Li filled p orbital Atomic number
66. First Ionization energy He N F H C O Be B Li Atomic number
67. He Ne Ne has a lower First Ionization energy N F IE than He Both are full, H C O Ne has more Be shielding B Greater Li distance Atomic number
68. He Ne Na has a lower First Ionization energy N F IE than Li Both are s1 H C O Na has more Be shielding B Greater Li distance Na Atomic number
69. First Ionization energy Atomic number
70. Driving Forces Full Energy Levels require lots of energy to remove their electrons. • Noble Gases have full orbitals. Atoms behave in ways to try and achieve a noble gas configuration.
71. 2nd Ionization Energy For elements that reach a filled or half-filled orbital by removing 2 electrons, 2nd IE is lower than expected. True for s2 Alkaline earth metals form 2+ ions.
72. 3rd IE Using the same logic s2p1 atoms have an low 3rd IE. Atoms in the aluminum family form 3+ ions. 2nd IE and 3rd IE are always higher than 1st IE!!!
73. Trends in Ionic Size: Cations Cations form by losing electrons. Cations are smaller than the atom they came from – not only do they lose electrons, they lose an entire energy level. Metals form cations. Cations of representative elements have the noble gas configuration before them.
74. Ionic size: Anions Anions form by gaining electrons. Anions are bigger than the atom they came from – have the same energy level, but a greater area the nuclear charge needs to cover Nonmetals form anions. Anions of representative elements have the noble gas configuration after them.
75. Configuration of Ions Ions always have noble gas configurations ( = a full outer level) Na atom is: 1s22s22p63s1 Forms a 1+ sodium ion: 1s22s22p6 Same configuration as neon. Metals form ions with the configuration of the noble gas before them - they lose electrons.
76. Configuration of Ions Non-metals form ions by gaining electrons to achieve noble gas configuration. They end up with the configuration of the noble gas after them.
77. Ion Group trends Each step down a group is adding Li1+ an energy level Na1+ K1+ Ions therefore get bigger as you go Rb1+ down, because of Cs1+ the additional energy level.
78. Ion Period Trends Across the period from left to right, the nuclear charge increases - so they get smaller. Notice the energy level changes between anions and cations. N3- O2- B 3+ F1- Li1+ Be2+ C4+
79. #3. Trends in Electronegativity Electronegativity is the tendency for an atom to attract electrons to itself when it is chemically combined with another element. They share the electron, but how equally do they share it? An element with a big electronegativity means it pulls the electron towards itself strongly!
80. Electronegativity Group Trend The further down a group, the farther the electron is away from the nucleus, plus the more electrons an atom has. Thus, more willing to share. Low electronegativity.
81. Electronegativity Period Trend Metals are at the left of the table. They let their electrons go easily Thus, low electronegativity At the right end are the nonmetals. They want more electrons. Try to take them away from others High electronegativity.
82. The arrows indicate the trend: Ionization energy and Electronegativity INCREASE in these directions
83. Atomic size and Ionic size increase in these directions:
+1-315-636-4485
+91 98151-74050 
