The Periodic Table

1. The Periodic Table Trends

2. Johann Dobereiner • Groups of 3 with similarities/trends – triads • Cl, Br, I: the properties of Br were between those of Cl and I • Limited to some groups, ineffective with others History

3. JAR Newlands • Law of Octaves • 1864 • Every eight elements the pattern repeats (like the musical scale) • He was viewed as a fool History

4. Original was arranged by MASS • Dmitri Mendeleev and J. Lothar Meyer (1869) • Mendeleev predicted existence of unknown elements (Ge, Sc, and Ga) and predicted their properties from the patterns he saw • Corrected the assumed atomic masses for elements (Be, U, In) • Dubbed the “Father of the Modern Periodic Table” over Meyer Modern Periodic Table

5. Ekasilicon…Germanium

6. Changed the table to be organized by atomic number (Z) • Then, it more closely followed trends and patterns Henry Mosley

7. Electron configuration trends are seen in the periodic table • Groups are based on 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) e- Configs and the Table

8. Blocks correspond to the orbital of the outermost electron Blocks and orbital shape (l)

10. Atomic number (not periodic) • Electron configuration • Atomic Radii • Ionization energy • Electron affinity • Electronegativity • Activity • Density Patterns and the Table

11. Mendeleev says: “The properties of the elements are a periodic function of their atomic masses” • Now we say: “When atoms are arranged by increasing atomic number, the physical and chemical properties show a repeating pattern” Periodic Law

12. Properties of elements are periodic functions of their atomic numbers • That’s why we call it the periodic table… Periodic Law

13. Atoms gain, lose, or share electrons in order to create a FULL outer shell • Typically, this is 8 electrons • The law can be used to predict several properties • H and He are exceptions: • H gains an electron to become H- (same e config as He) • H may want to go to no electrons which is “full” although it’s actually empty Octet Rule

14. Half the distance between adjacent nuclei • ½ (2r) = atomic radius Atomic Radii

15. INCREASES down a group (because n increases) • DECREASES across a period • As electrons are added, a proton is also added so the nucleus becomes more positively charged (electrons have same negative charge) • Results in each electron being MORE attracted to the INCREASINGLY more positive nucleus and being PULLED in closer • Think of a magnet Atomic Radii

17. Cations (positive) • Smaller than the neutral atom • The electrons have less repulsion and pull in closer to the nucleus • Anions (negative) • Larger than the neutral atom • More electrons = more repulsion = larger electron cloud Ionic Radii

18. Ionic Radii

19. This is the amount of energy needed to remove an electron from an atom • More specifically, an isolated atom of the element in the gas phase • Usually measured in kJ/mol Al(g)Al(g)+ + e- I1 = 580 kJ/mol Al(g)+ Al(g)+2 + e- I2 = 1815 kJ/mol Ionization Energy

20. First IE: energy needed to remove the first electron from an element • Second IE: energy needed to remove the second electron from an element Ionization Energy

21. There are more…3rd, 4th, 5th and so on until you cannot pull any more electrons off (successive IE’s) • It takes more energy to remove successive electrons than to remove the first • Because… 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 • 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 Ionization Energy

22. Electrons want to hang around the atom so it takes energy to remove electrons • In general • The smaller the 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 positively charged nucleus • Also, less repulsion between electrons Ionization Energy

24. Ionization Energy

25. Inner core electrons are those electrons from the previous Noble Gas (abbreviated electron configurations) • Valence electrons are the electrons that are on the exterior of the atom • These are the electrons that are responsible for properties or behavior of the element Remember

26. Higher than the first • Because there are 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 • Increasingly larger jumps as each electron is removed • One jump is much larger than the others, because once the inner core configuration is reached, electrons are removed from the inner core, taking a lot more energy Successive IEs

27. Successive IEs

29. The ability of an atom to attract electrons in a bond • Some atoms share electrons easily, others do not • The ability to share is rated from 0 to 4 • Elements with 0 electronegativity share easily • Elements with a high electronegativity do not share electrons well Electronegativity

30. If the element normally forms a cation, it has a low electronegativity • If the element normally forms an anion, it has a high electronegativity • The smaller it is, the higher the electronegativity • The larger it is, the lower the electronegativity • Noble gases (take no charge) have NO electronegativity values Electronegativity Trends

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

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

34. 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 Metals Reactivity Trend

35. In groups, reactivity of nonmetals decreases when atomic number increases • Because the electronegativity decreases • Relate to size (it increases) • In periods, reactivity of nonmetals increases with atomic number • Because the electronegativity increases • Relate to size (radii decreases) Nonmetals Reactivity Trend

36. 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 • Decreases down a group • Decreases across a period • Density of liquids • Measured in g/mL • Density of Hg is greater than that of Br2 Density

37. The amount of energy released when an electron is added to a neutral atom Electron Affinity