The Periodic Table

1. 0 The Periodic Table Beyond protons, neutrons, and electrons

2. 0 Early PT Folks • Johann Dobereiner • Triads- groups of 3 with similarities/ trends • Cl, Br, I – the properties of Br were intermediate to those of Cl and I • Limited to some groups, not effective with others • JAR Newlands (1864) Law of Octaves • Every eight elements the pattern repeats itself, similar to a musical scale repeating every 8 notes • Not generally well received; people thought him a fool

3. 0 The Modern Periodic Table • The original PT was arranged by mass • By Dmitri Mendeleev and J Lothar Meyer in 1869 • Mendeleev predicted the existence of unknown elements (which turned out to be Ge, Sc, and Ga), and predicted their properties from the patterns he saw • Mendeleev corrected the assumed atomic masses for elements (In, Be, U) • These are reasons why he is credited with the first periodic table and is dubbed “The Father of the Modern Periodic Table” over Meyer

4. 0 Ekasilicon

5. 0 Changes…. • Henry Mosley changed the table to be organized by atomic number (Z) instead; it then more closely followed trends/ patterns

6. 0 e- configuration and the PT • PT also shows trends in electron configuration • Groups are based upon electron configuration • Alkali metals are #s1 (# is period) • Alkaline earth metals are #s2 (# is period) • Halogens #p5 (# is period) • Noble gases #p6 (# is period) • Transition metals d block (# is period -1) • Inner transition metals are f block (# is period -2)

7. 0 Blocks and l* ** orbital shape The blocks you already know correspond to the orbital of the last (outermost) e- , or valence e-s occupied

8. 0 Patterns (Periods) and the PT • We see patterns for many things, including • Atomic number *(not a periodic pattern, but a pattern) • Electron configuration • Atomic radii • Ionization energy • Electron affinity • Electronegativity • Activity • Density

9. 0 The Periodic Law • Mendeleev says "The properties of the elements are a periodic function of their atomic masses" • We now say: “When atoms are arranged by increasing atomic number, the physical and chemical properties show a (repeating) pattern”

10. 0 Periodic… • Summed up: Properties of elements are periodic functions of their atomic numbers. • Hence, we call the table of elements the PERIODIC table (go figure)

11. 0 Octet Rule • “Atoms gain, lose, or share electrons in order to create a full outer shell” • This is typically going to be eight electrons • H and He are exceptions; wanting to fill the 1s orbital • H gains an electron to become H- , with the same electron configuration as He • H may want to go to no electrons, which is considered “full” even though it is empty • H+ and He+2 would have no electrons left • The law can be used to predict several properties

12. 0 Atomic Radii • Half the distance between adjacent nuclei • ½ (2R)= atomic radius

13. 0 Atomic Radii • The radius increases as you go down a group • This is because n increases • The radius decreases as you go across a period (Yes, this is counterintuitive) • Due to the fact that you add e- as you add p+, so the nucleus is more positively charged, and each electron has the same negative charge • Results in each electron being more attracted to the (increasingly) more positive nucleus, and being pulled in closer • Sort of like making a magnet more powerful- it will decrease the distance where it will pull objects towards it

15. 0 Ionic Radii • Cations (+) • Smaller than the neutral atom • The electrons have less repulsion, and pull in closer to the nucleus • Anions (-) • Larger than the neutral atom • More electrons = more repulsion = larger electron cloud

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17. 0 Ionization Energy (Heretofore called IE) • IE is the amount of energy needed to remove an electron from an atom • (specifically, an isolated atom of the element in the gas phase) • Measure in kJ/ mol Al(g)Al(g)+ + e- I1 = 580 kJ/mol Al(g)+ Al(g)+2 + e- I2 = 1815 kJ/mol

18. 0 IE, continued • The Energy needed to remove the first electron from an element is the 1st IE • The Energy needed to remove the second electron is known as the 2nd IE

19. 0 Successive IE • There are also 3rd, 4th, 5th , and so on IEs (which are successive IEs), until you can’t pull any more off • It takes more energy to remove successive electrons than to remove the first • Due to the fact that there are then more protons than electrons, and the stronger positive charge will then act on the remaining electrons to hold them to the atom • (Remember that the charge on the nucleus increases while the charge on each electron remains the same, causing more pull by the nucleus on each individual electron)

20. 0 Why IE? • Since electrons (-) want to hang around the atom (due to the + protons in the nucleus pulling on them), it takes energy to remove electrons • In general • The smaller that atom, the more energy it takes to remove an electron • Because the electron is closer to the nucleus than in a larger atom • The fewer electrons that atom possess, the harder it is to remove an electron • Because it will hang on to them tighter as they are closer to the + charged nucleus; • also, the less repulsion between electrons

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22. 0 1st IE

23. 0 Stuff to keep in mind… • Remember (from coming up with the abbreviated electron configurations) that: • Inner core electrons are those electrons from previous Noble Gas • Valence electrons are the electrons that are on the exterior of an atom • These are the electrons that are responsible for the behavior (properties) of the element

24. Successive IEs • Are higher than the first • Due to the fact that there is going to be more protons than electrons at that point, resulting in a stronger attraction on the remaining electrons than there was in the first place • Basically increasingly larger jumps as each electron is removed • One jump is usuallymuch larger than the others, because once the inner core configuration is reached, electrons are removed from the inner core, taking a lot more energy • Much bigger difference between positive nucleus and negative electron

25. Successive IEs

26. IE and the PT

27. Electronegativity (Eneg) • The ability of an atom to attract electrons in a bond • Some atoms share electrons easily, others are electron hogs • The ability to share is rated (usually) from 0 to 4 • Elements with 0 Eneg share easily • Elements with a high (close to 4) Eneg don’t share e- well

28. Electronegativity Trends • If it normally goes +, it has a low Eneg • If it normally goes -, it is has a high Eneg • The smaller it is, the higher the Eneg • The larger it is, the lower the Eneg • Noble gases, which normally take no charge, we say have no Eneg values

29. Electronegativity Trends

30. Metallic character • Metallic character is acting like a metal (conductive, shiny, malleable,etc) • All elements possess from very low to very high metallic character. • The scale is from Fr to F. • Fr has the most metallic character and F has the least. • In groups,  metallic character increases with atomic number because each successive element gets closest to Fr. • In periods,  metallic character decreases when atomic number increases because each successive element gets closest to F.

31. Reactivity • The nature (metal, non-metal, semi-metal) makes a difference in how an element’s chemical reactivity • The trends are characterized by their nature

32. Metals reactivity trend • In groups, reactivity of metals increases with atomic number because the ionization energy decreases. • In periods, reactivity of metals decreases when atomic number increases because the ionization energy increases.

33. Nonmetals reactivity trend • In groups, reactivity of non-metals decreases when atomic number increases • because the electronegativity decreases • Relate to size- it increases. • In periods, reactivity of non-metals  increases with atomic number • because the electronegativity increases. • Relate to size- radii decreases • Remember, the radii would have an effect on this

34. Density: in general • Density of solids is greatest • Measured in g/cm3 • Highest in center of table (d- block) • Density of gases • Measured in g/L at STP (1atm , 0°C) • Decreases as you go down a group • Decreases as you go across the table • Density of liquids • Measured in g/mL • Density of Hg is greater than that of Br2

35. Density

36. Density v atomic number

37. To sum it up…

38. Summing it up (again)

39. Summary chart again