Periodic Table – Organizing the Elements

1. Periodic Table – Organizing the Elements Chapter 5.4 & Chapter 14

2. Dmitri Mendeleev • About 70 elements had been found by the mid 1800’s • Mendeleev was the first to organize them in a systematic way

3. He listed the elements in order of increasing atomic mass • Arranged the elements in columns so those with similar properties were side by side

4. He left blank spaces where nothing fit • He predicted the physical properties of the missing elements • He was mostly correct

5. Henry Moseley • Moseley determined the atomic # of the elements and arranged the table by atomic number instead of atomic mass

6. The modern periodic table is arranged by atomic number • The periodic table has atomic # increasing from left to right & top to bottom

7. Periodic Law • The horizontal rows on the periodic table are called periods • Properties change as you move across a period

8. The properties repeat when you move from one period to the next • Periodic Law: there is a periodic repetition of the chemical & physical properties of the elements

9. Groups • Each vertical column is called a group or family • Elements in the same group have similar properties

10. Groups have a number and a letter (pg 124) • The group with Li, Na, K etc is called Group 1A • Group 1A elements are also called the alkali metals

11. All group A elements are called the representative elements • They exhibit a wide range of physical & chemical properties

12. Elements on the left side of the periodic table (except for hydrogen) are metals • Group 2A are the Alkaline Earth Metals

13. Group B elements are the transition & inner-transition metals • Gold & silver are transition metals • Uranium is an inner-transition metal

14. The upper right hand corner of the table has the non-metals • Some are gases, some are solids & some are liquids at room temperature

15. Bromine is a liquid, Oxygen is a gas and sulfur is a solid • Group 7A are called the Halogens (F, Cl, Br, I) • Group 0 (8A) are the noble gases (He, Ne, etc)

16. Metals have a shiny appearance (luster) & are good conductors of heat & electricity, most are solids • Nonmetals do not have luster & are poor conductors

17. Elements bordering the step-line are called metalloids or semi-metals • Si & Ge are metalloids • They have properties in-between metals & nonmetals

18. The elements can also be classified by their electron configuration • Electrons play the most important part in determining the properties of elements

19. Write the electron configurations for the Alkali Metals • What similarities do you see? • The Halogens? • The Noble Gases?

20. The noble gases have their outermost s & p sublevels filled completely • The Representative Elements have their outermost s & p sublevels partially filled

21. The Transition Metals – their outermost s & nearby d sublevels contain electrons • The Inner Transition Metals – their outermost s & nearby f sublevels contain electrons

22. The Table can be broken up into blocks - tell you the outermost sublevels that are filled • s block, p block, d block & f block • Where are they?

23. Each period on the Table corresponds to a principle energy level being filled • # electrons can be determined by counting left to right • d block is one less than the period, f block 2 less

24. The electron configuration can be determined for most of the elements this way

25. “S” block • Groups 1 & 2 • Electron Configuration ends in an S Sub-level. • Highest energy level is equal to the period number of the element. • i.e. Calcium’s (in the 4th period) electron configuration ends in 4s.

26. “P” block • Groups 13 thru 18 • Electron Configuration ends in a P Sub-level. • Highest energy level is equal to the period number of the element. • i.e. Silicon’s (in the 3rd period) electron configuration ends in 3p.

27. “D” block • Groups 3 thru 12 • Electron Configuration ends in a D Sub-level. • Highest energy level one less than the period number of the element. • i.e. Silver’s (in the 5th period) electron configuration ends in 4d.

28. “F” block • “Inner Transition Metals” • Electron Configuration ends in an F Sub-level. • Highest energy level two less than the period number of the element. • i.e. Uranium’s (in the 7th period) electron configuration ends in 5f.

29. Regardless of the “Block,” the number of electrons in the highest sub-level is equal to the element’s column number within its block. • Ex: Nitrogen is in the 3rd column of the p block and its configuration ends in p3. • Ex: Iron is in the 6th column of the d block and it ends in d6.

30. Examples: • Determine the last term in the electron configurations of the following elements: • Chlorine: • 3p5 • Potassium: • 4s1 • Mercury: • 5d10

31. Stable electron configurations 1. Ending in a full p sublevel or 1s • Noble gases are the most stable 2. Ending in a full sub-level other than a p or the 1s. • Magnesium (3s2) is more stable than Sodium (3s1).

32. 3.Ending in a half-filled multi- orbital (p, d or f) sub-level. • Nitrogen (2p3) is more stable than Carbon (2p2) or Oxygen (2p4). • Stability due to un-paired electrons in orbitals having same “Spin.”

33. Periodic Trends • An element’s placement in the periodic table determines characteristics like the size of the atom, its ability to attract electrons and the stability of its electron configuration.

34. Atomic Radius • Size of atoms of each element: • How will the size of atoms change as we proceed down a group? • i.e. Compare the sizes of Li and Na. • From Li to Na, we add an entire energy level, therefore the size increases.

35. How will the size of atoms change as we proceed across a period? • Compare C, N and O. Which is largest? • Oxygen has the most electrons. • However, it also has the most protons. • The outermost electrons of Oxygen are in the same sub-level as C and N.

36. Oxygen’s greater nuclear charge attracts the electrons, causing the atom to contract! • Oxygen is the smallest of the three, Carbon is the largest. • Atomic Radius decreases as we go across a period from left to right and up a group.

37. Examples • Rank the following sets in order of decreasing Radius. • S, Cr, Se, Sr, Ne • Sr, Cr, Se, S, Ne • Fe, N, Ba, Ag, Be • Ba, Ag, Fe, Be, N

38. Ionic Size vs. Atomic Size • When an atom becomes an ion, it will either get smaller or larger • Metals lose electrons and will get smaller (stronger pull from the nucleus) • Nonmetals gain electrons and will get bigger (more e- to repel one another)

39. Ionic Size vs. Atomic Size • Which is bigger? • Na or Na+ • Cl or Cl- • O or O-2 • Mg or Mg+2

40. Ionization Energy • Amount of energy required to remove a valence electron from an atom. • The more stable an element is, the harder it will be (more energy is required) to remove an electron. • Some elements become more stable by losing an electron so they lose electrons easily (less energy needed).

41. How does ionization energy vary within a group (compare Li and Na)? • The electron to be removed from Na is further from the nucleus than Lithium’s electron. • Sodium’s electron is held more loosely and therefore easier (less energy) to remove.

42. How does ionization energy vary across a period? (Compare elements in 3rd period) • Sodium attains a Noble Gas configuration by losing an electron, so little energy is required. • Magnesium is somewhat stable due to a full 3s sub-level, so more energy is needed.

43. Argon is a Noble Gas. Due to its stability, it is very difficult (much energy needed) to remove an electron. • Chlorine has no stability in its configuration, so it is easier to remove an electron. • Ionization energy increases across a period and up a group.

44. Examples • Rank the following sets in order of decreasing Ionization Energy. • K, Zn, Cs, Ar, P • Ar, P, Zn, K, Cs • C, He, Ag, Pt, Sn • He, C, Sn, Ag, Pt

45. 2nd Ionization Energy • The 2nd ionization energy is the amount of energy required to remove the second electron on the outside of an atom • Sometimes it is larger than the 1st ionization NRG, sometimes, it is smaller

46. 2nd Ionization Energy • For elements like Na and the alkali metals, the 2nd ionization NRG is much higher than the 1st • WHY??? • Na loses 1 e- and becomes like a noble gas. • Losing the 2nd would be counter-productive and will not happen easily!

47. 2nd Ionization Energy • For elements like Mg and the other alkaline earth metals, the 2nd ionization NRG is lower than the 1st • Losing 1 e- is relatively difficult because of the s2 configuration (somewhat stable) but losing the next e- is super easy

48. 2nd Ionization Energy • How do you think the 1st ionization NRG and the 2nd ionization NRG compare for the halogens? • The Noble gases? • WHY???

49. Electronegativity • Describes an element’s attraction for an electron in a covalent bond. • Elements that need electrons to complete an energy-level will have a high electronegativity. • Elements that want to lose electrons have low electronegativities.

50. How does Electronegativity vary within a group? (compare F and Cl) • Both elements need an electron to complete a p sub-level. • Fluorine’s p sub-level is closer to its nucleus, so it has a greater magnetic attraction for a free electron. • F has a higher electronegativity!