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!

51. How does electronegativity vary across a period? (period 2) Fluorine benefits the most by gaining an electron, so it has the highest electronegativity. Lithium, which wants to lose an electron has very little attraction for an additional electron.

52. Carbon can gain electrons but sometimes loses them as well, so its electronegativity is between F and Li. Noble Gases have no Electronegativity! Electronegativity increases across a period and up a group (Noble Gases omitted).

53. Examples Rank the following sets in order of decreasing Electronegativity: Cu, F, Mn, Sr, Si F, Si, Cu, Mn, Sr Al, Ca, S, Cl, Fe Cl, S, Al, Fe, Ca