1. Chapter 6 “The Periodic Table” Edited from “Pre-AP Chemistry” By S.L. Cotton Posted on website!

2. Section 6.1 Organizing the Elements u A few elements, such as gold and copper, have been known for thousands of years - since ancient times u Yet, only about 13 had been identified by the year 1700. u As more were discovered, chemists realized they needed a way to organize the elements. 2

3. Mendeleev’s Periodic Table u Chemists used the properties of elements to sort them into groups. u By the mid-1800s, about 70 elements were found or known to exist u Dmitri Mendeleev – a Russian chemist and teacher u Arranged elements in order of increasing atomic mass (incorrectly!) u Created the first “Periodic Table” 3

4. Mendeleev u He left blanks for predicted yet undiscovered elements where masses or reactivities were missing. Later the ‘blank’ elements were discovered, and his ‘periodic’ order was proven correct! u But, there were a few problems: Such as Co and Ni; Ar and K; Te and I which didn’t behave like the others in their group (based on mass) 4

5. A Better Arrangement u In 1913, Henry Moseley – British physicist, discovered the proton and re-arranged elements according to increasing atomic number u The arrangement used today u The same symbol, atomic number & mass are basic items used today 5

6. 118 6 Moseley’s Table u Horizontal rows = periods There are 7 periods u Vertical column = group (or family since they have similar properties) There are 18 groups, 8A groups (valence e-) u Identified by number and or letter (IA, 1)

7. Spiral Periodic Table Early Table: Spiral Periodic Table 7

8. The Periodic Law: u When elements are arranged in order of increasing atomic number, there is a ‘periodic’ repetition of their physical and chemical properties. 1234321 (ratios) u These similar physical & chemical properties are really due to the group/family having identical valence electron shell. u EX halogens 7e-, Noble gases, 8 e- 8

9. Areas of the periodic table u Three main classes of elements are: 1) metals, 2) nonmetals, and 3) metalloids (aka semi-metals) 1) Metals: electrical conductors, have luster (shine), ductile (wire), malleable, solids 2) Nonmetals: generally brittle and non-lustrous, poor conductors of heat and electricity, some liq, gas. 9

10. Areas of the periodic table u Some nonmetals are gases (O, N, Cl); some are brittle solids (S); one is a fuming dark red liquid (Br) u Notice the heavy, stair-step line? 3) Metalloids: border the line-2 sides Can have properties are of both metals and nonmetals depending on conditions 10 “Stair” Metal/Nonmetal divider S block 2 e max d block 10e max p block 6 e max f block 14 e max

11. Squares in the Periodic Table The periodic table displays the symbols and names of the elements, along with information about the structure of their atoms: 11

12. Groups ‘Families’ of elements u These “families” have similar properties. This is due to their identical valence electron configuration! u Group I A– alkali metals (1 valence e-) Forms a “base” (or alkali) when reacting with water (not just dissolved!) u Group II A – alkaline earth metals (2e-) Also form bases with water; do not dissolve well, hence “earth metals” u Group IIIA (13) – Boron Group 12

13. Groups ‘Families’ of elements u Group IV A (14) – Carbon Group 4 valence e- u Group V A (15) Nitrogen Group 5valence e- u Group VI A (16) Oxygen Group u 6 valence e- u Group 17 (7A)– halogens (7e-_ Means “salt-forming” NaCl, KI, etc. 13

14. Group 1(1A) are the alkali metals (but NOT H) Group 2(2A) are the alkaline earth metals Group 17(7A) is called the halogens H 14

15. Electron Configurations in Groups 1) Noble gases are the elements in Group 18 (also called Group 8A 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 15

16. Why do some chemicals react? u Octet Rule: Full S 2e -& P 6e- energy levels are most stable 8! and require lots of energy to remove their electrons (they don’t follow trends) Noble Gases have full orbitals. u Elements react in ways to try and achieve a stable noble gas electron configuration to get the ‘octet’ in their valence 16

17. u Group 18 (8A) are the noble gases 17

18. Electron Configurations in Groups Transition metals are in the roman numeral “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 18

19. Electron Configurations in Groups 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 19

20. 3-12 aka group 3B-12B are called the transition elements u These are called the inner transition elements aka u Lanthanides :within p6 u Actinides within p7 u Together known as the Rare Earth Elements 20

21. 1s11s1 1s 2 2s 1 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6 s 2 4f 14 5d 10 6p 6 7s 1 H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87 Do you notice any similarity in these e- configurations of the alkali metals? Electron Configurations & Orbitals Number is shell, Letter is orbital, exponent is electron amount. 21

22. He 2 Ne 10 Ar 18 Kr 36 Xe 54 Rn 86 1s21s2 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 2 3p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 Do you notice any similarity in the electron configurations of the noble gases? 22

23. u Each row (or period) is the energy level for s and p orbitals. 12345671234567 Period Number 23

24. Periodicity ( Repeating Patterns ) on the Table 1. Atomic Radius How “big” the atom is from edge of the valence to center. (mainly determined by # of “shells” and pull from protons in nucleus) 2. Electronegativity: How much an atom desires or pulls electrons toward it. How much it wants to bond e-. (depends how much protons pull.) u 3. Ionization Energy (1 st ) u How much energy in eV or J needed to remove one valence electron (depends on pull from + nucleus) 24

25. ALL Periodic Table Trends u Influenced by three factors: 1. Energy Level Higher energy levels ‘shells’ are further away from the nucleus. 2. Charge on nucleus (# protons) More charge pulls electrons in closer. (+ and – attract each other) u 3. Shielding effect of each electron shells can block pull of protons in nucleus 25 Periodicity Activity!

26. #1. Atomic Size/Radius - Group trends u Measured in A, angstrom 10 -10 m u As we go down a group... u each atom has another ‘shell or energy level, u so the atoms get bigger. H Li Na K Rb 26

27. #1. Atomic Size - Period Trends u Going from left to right across a period, the size gets smaller. u Electrons are in the same energy level. u But, there is more positive nuclear charge. u Outermost electrons are pulled closer. NaMgAlSiPSClAr 27

28. Atomic size and Ionic size increase in these directions: Increases 28

29. Ions u Metals tend to LOSE electrons, from their outer energy level because this will most easily leave the full outer shell beneath (losing 1 is easier than gaining 7 to get octet) EX Sodium will: Lose 1 electron there are now more protons (11) than electrons (10), and thus a positively charged particle is formed = “cation” The charge is written as an exponent followed by a plus sign: Na 1+ Now named a “sodium ion” 29

30. Ions u Nonmetals tend to GAIN one or more electrons EX. Chlorine tends to: gain one electron to get 8 Protons (17) no longer equals the electrons (18), so a charge of -1 Cl 1- is re-named a “chloride ion” Negative ions are called “anions” 30

31. #2 Ionization Energy: is the energy needed to pull off electrons. It increases across the periods due to increasing proton, increasing attraction for electrons. (This is like the opposite of Atomic Size trend in a period) 31 First Ionization Energy decreases downward because the additional electron shells are farther from positive nucleus attraction and are shielded by each shell

32. #3. Trends in Electronegativity u Electronegativity is the tendency for an atom to attract electrons to itself when it is chemically combined with another element. u An element with a big electronegativity means it pulls the electron towards itself strongly! 32

33. Electronegativity Group Trend u The further down a group, the farther the electron is away from the nucleus, plus the more electrons an atom has. u Thus, more willing to share and easier to steal. u Low electronegativity. 33 H Li Na K Rb

34. The arrows indicate the trend: Electronegativity INCREASES in these directions toward Fluorine Increases u Metals are at the left of the table. u They let their electrons go easily u Thus, low electronegativity u At the right end are the nonmetals. u Nonmetals want more electrons. u Try to take them away from others so u High electronegativity. 34