Unit 8 Periodic Trends

1. Unit 8 Periodic Trends Textbook chapter 5 http://www.youtube.com/user/Hermcast

2. 5.1 HISTORY OF THE PERIODIC TABLE • Brainstorm some possible ways to organize the about 70 elements known by the mid-1800s.

3. 5.1 HISTORY OF THE PERIODIC TABLE • Brainstorm some possible ways to organize the about 70 elements known by the mid-1800s. • Alphabetic • Physical states

4. 1817 Johann Döbereiner’s. Organized elements into ______triads_____, similar to atomic __family_ or groups.

5. http://en.wikipedia.org/wiki/Johann_Wolfgang_D%C3%B6bereiner • For example, the average atomic mass of lithium and potassium was close to the atomic mass of sodium. A similar pattern was found with calcium, strontium, and barium, with sulphur, selenium, and tellurium, and also with chlorine, bromine, and iodine. • http://www.merriam-webster.com/cgi-bin/audio-medlineplus.pl?bixdob01=Dobereiner

6. 1863 John Newlands. Arranged elements in order of increasing atomic masses. Noticed a pattern formed by every 8th element, called it the “Law of octaves”. Remember, at this time the entire ___noble gas__ family was unknown.

7. 1869 Dmitri Mendeleev, know as the “The father of the periodic table__”. Agreed with Newlands’ arrangement, but added transitions, giving the periodic table its current shape.

9. Left blank spots on the chart for adding in as yet undiscovered elements. • Made predictions about the chemical and physical properties of these elements. His most famous example of this is regarding Ekasilicon, or modern day Germanium.

11. Mendeleev’s Periodic Law states: • “The physical and chemical properties of elements are periodic functions of their atomic mass."

12. The main value of the periodic table is the ability to predict the chemical properties of an element based on its location on the table. It should be noted that the properties vary differently when moving vertically along the columns of the table than when moving horizontally along the rows.

13. Lothar Meyer 1869. Worked on the periodic table independent of Mendeleev and came to the same conclusions. Didn’t add predictions and so didn’t get the credit.

14. 1911 Henry Moseley Noticed 3 pairs of elements out of place. They are:__________, _____________, and _____________. Restated the Periodic Law as :

15. 1911 Henry Moseley Noticed 3 pairs of elements out of place. They are: Te + I, Co + Ni, and Ar + K. Restated the Periodic Law as : • “The physical and chemical properties of elements are periodic functions of their atomic number."

16. 5.2 Electron Configurations and the Periodic Table • Periods and Blocks of the Periodic Table • “s” Block = groups # __1 ______ and ______2_______ •   s = sharp • “p” Block = groups # __13__________ thru ____18______ •   p = principal • “d” Block = groups # ____3_______ thru ______12_____ •   d = diffuse • “f” Block = groups # Lanthanide and Actinide •   f = fundamental

17. 5.3 Electron Configuration and Periodic Properties • Atomic radii is defined as: The distance between nuclei of 2 atoms.

18. Measured by what technique? Xray diffraction

19. Need a blank sheet of paper • Draw: • Li Be B C N O F Ne • Na • K • Rb See hermcast for an added explanation

20. TRENDS for atomic size: • Group: Atomic size ___________ as you move up a group. • Period: Atomic size________ as you move left to right across a period.

21. On back of blank sheet of paper • Atomic size: Periodic Table See hermcast for an added explanation

22. Shielding effect explanation: For elements with several electrons arranged over several shells (energy levels), inner electrons block (or shield) outer or valance electrons from the positive charge of the nucleus.

23. Shielding effect explanation: For elements with several electrons arranged over several shells (energy levels), inner electrons block (or shield) outer or valance electrons from the positive charge of the nucleus. This causes a decrease in charge attractions, making the outer electrons easier to lose.

25. The following three slides are not part of the notes but may help

26. 1. Atomic Radii (p150) • periodic trend (left to right)– decreases • due to increasing charge of the nucleus • Why does charge increase? • group trend (top to bottom)– increases • due to shielding by other electrons • What is shielding?

27. 1. Atomic Radii (p150) • periodic trend (left to right)– decreases • due to increasing charge of the nucleus • Why does charge increase? • As you move from left to right the number of protons increase • Since there are more protons in the nucleus and they all have a positive charge they pull the electrons moving around them closer to the nucleus this makes the elements on the right smaller than the elements on the left.

28. 1. Atomic Radii (p150) • group trend (top to bottom)– increases • due to shielding by other electrons • What is shielding? • As you move down a group the size of atoms increase because you are adding energy levels • Energy levels are the areas electrons are moving around the nucleus. • Since electrons don’t want to be near each other since the are the same charge there is a large gap between energy levels • This makes the atoms larger as you move down

29. Periodic Trends in Ionization EnergyA + Energy  A+ + e- • Ionization Energy is the amount of energy needed to lose an electron. This causes the atomic to become an ion.

30. Periodic Trends in Ionization EnergyA + Energy  A+ + e- • Ionization Energy is the amount of energy needed to lose an electron. This causes the atomic to become an ion. • Metals tend to form cations, low ionization energy value.

31. Nonmetals hate forming cations, instead prefer to form anions. They have a high ionization energy.

32. First ionization Energy is the amount needed to remove the _first___ electron. • Second ionization Energy is the amount needed to remove the __second__ electron. • Third ionization Energy is the amount needed to remove the __third__ electron.

33. TRENDS in ionization energy: • Group: Tends to increase as you move up a group. • Period: Tends to increase as you move left to right across a period. • Highest ionization energy? • Lowest?

34. On back of blank sheet of paper • Ionization energy: Periodic Table See hermcast for an added explanation

35. The following four slides are not part of the notes but may help

36. 2. Ionization energy (p154) • element + energy  Element+ + e- • def- energy required to remove one electron from a neutral atom (first ionization energy) • periodic trend – increases • due to the electrons become closer to the nucleus • group trends – decreases • due to shielding

37. 2. Ionization energy (p154) • element + energy  Element+ + e- • def- energy required to remove one electron from a neutral atom (first ionization energy) • periodic trend – increases • due to the electrons become closer to the nucleus • This holds the electrons more strongly than if the electrons were further apart • Since they are held on more strongly it takes more energy to pull them off and remove them • Therefore there is a larger number associated with energy

38. 2. Ionization energy (p154) • H + energy  H+ + e- • group trends – decreases • due to shielding • It’s easier to pull electrons off since they are further away from the nucleus. • They are further away since as you move down there are move energy levels

39. second ionization energy – removal of a second electron (this requires more energy) • Why does it require more energy? • Once you remove one electron now you have a positively charged ion • Na  Na+ +e- • Na+ is already positively charge trying to make it Na+2 requires A LOT of energy • once you remove enough electrons to achieve an electron configuration of a noble gas it is very difficult to remove an additional electron (requires a LOT of energy)

40. Electron Affinity A + e- A- + Energy • Electron affinity: a measure of the attraction an atom has for gaining electrons. • An affinity is an attraction for something.

41. Metals have _low_ attraction for gaining electrons, low electron affinity values.

42. Nonmetals have _high__ attraction for gaining electrons, high electron affinity values.

43. Feel free to LOL • Mr. Stevenson has ____________attraction for staples at a 45 degree angle ,  _________ staple affinity values.

44. TRENDS for electron affinity • Group: Tends to increase as you move up the group. • Period: Tends to increase as you move left to right across a period.

46. On back of blank sheet of paper • Electron affinity size: Periodic Table See hermcast for an added explanation

47. The following slide is not part of the notes but may help

48. electron affinity (p158) • periodic trend (left to right) – generally increases • nonmetals want to gain e- to have 8 valance electrons • Since size decreases they have the ability to gain electrons and they want to gain electrons to get 8 in their outer energy level since that is the most stable. • group trend – generally decreases • higher energy levels makes it more difficult for the extra electrons to be held to the atom • Shielding effect

49. Periodic Trends in Ionic Size • This is always a comparison between an ion and an atom of the same element.

50. Metals lose electrons increasing the nuclear charge, causing the ion to be __smaller_ than the atom.