2. Periodic Trends Elemental Properties and Patterns

3. The Periodic Law Dimitri Mendeleev was the first scientist to publish an organized periodic table of the known elements. He was perpetually in trouble with the Russian government and the Russian Orthodox Church, but he was brilliant never-the-less.

4. The Periodic Law Mendeleev even went out on a limb and predicted the properties of 2 at the time undiscovered elements. He was very accurate in his predictions, which led the world to accept his ideas about periodicity and a logical periodic table.

5. The Periodic Law Mendeleev understood the ‘Periodic Law’ which states: When arranged by increasing atomic number, the chemical elements display a regular and repeating pattern of chemical and physical properties.

6. The Periodic Law Atoms with similar properties appear in groups or families (vertical columns) on the periodic table. They are similar because they all have the same number of valence (outer shell) electrons, which governs their chemical behavior.

7. A Different Type of Grouping Besides the 4 blocks of the table, there is another way of classifying element: Metals Nonmetals Metalloids or Semi-metals. The following slide shows where each group is found.

8. Metals, Nonmetals, Metalloids

9. There is a zig-zag or staircase line that divides the table. Metals are on the left of the line, in blue. Nonmetals are on the right of the line, in orange.

10. Metals, Nonmetals, Metalloids Elements that border the stair case, shown in purple are the metalloids or semi- metals. There is one important exception. Aluminum is more metallic than not.

11. Metallic Character This is simple a relative measure of how easily atoms lose or give up electrons.

12. Metals, Nonmetals, Metalloids How can you identify a metal? What are its properties? What about the less common nonmetals? What are their properties? And what the heck is a metalloid?

13. Metals Metals are lustrous (shiny), malleable, ductile, and are good conductors of heat and electricity. They are mostly solids at room temp. What is one exception?

14. Nonmetals Nonmetals are the opposite. They are dull, brittle, nonconductors (insulators). Some are solid, but many are gases, and Bromine is a liquid.

15. Metalloids Metalloids, aka semi-metals are just that. They have characteristics of both metals and nonmetals. They are shiny but brittle. And they are semiconductors. What is our most important semiconductor?

16. Periodic Trends There are several important atomic characteristics that show predictable trends that you should know. The first and most important is atomic radius. Radius is the distance from the center of the nucleus to the “edge” of the electron cloud.

17. Atomic Radius Since a cloud’s edge is difficult to define, scientists use define covalent radius, or half the distance between the nuclei of 2 bonded atoms. Atomic radii are usually measured in picometers (pm) or angstroms (Å). An angstrom is 1 x 10 -10 m.

18. Covalent Radius Two Br atoms bonded together are 2.86 angstroms apart. So, the radius of each atom is 1.43 Å. 2.86 Å 1.43 Å

19. Atomic Radius The trend for atomic radius in a vertical column is to go from smaller at the top to larger at the bottom of the family. Why? With each step down the family, we add an entirely new PEL to the electron cloud, making the atoms larger with each step.

20. Atomic Radius The trend across a horizontal period is less obvious. What happens to atomic structure as we step from left to right? Each step adds a proton and an electron (and 1 or 2 neutrons). Electrons are added to existing PELs or sublevels.

21. Atomic Radius The effect is that the more positive nucleus has a greater pull on the electron cloud. The nucleus is more positive and the electron cloud is more negative. The increased attraction pulls the cloud in, making atoms smaller as we move from left to right across a period.

22. Effective Nuclear Charge What keeps electrons from simply flying off into space? Effective nuclear charge is the pull that an electron “feels” from the nucleus. The closer an electron is to the nucleus, the more pull it feels. As effective nuclear charge increases, the electron cloud is pulled in tighter.

23. Atomic Radius The overall trend in atomic radius looks like this.

24. Cation Formation 11p+ Na atom 1 valence electron Valence e- lost in ion formation Effective nuclear charge on remaining electrons increases. Remaining e- are pulled in closer to the nucleus. Ionic size decreases. Result: a smaller sodium cation, Na +

25. Anion Formation 17p+ Chlorine atom with 7 valence e- One e- is added to the outer shell. Effective nuclear charge is reduced and the e- cloud expands. A chloride ion is produced. It is larger than the original atom.

26. © 2009, Prentice-Hall, Inc. Sizes of Ions Cations are smaller than their parent atoms. The outermost electron is removed and repulsions between electrons are reduced.

27. © 2009, Prentice-Hall, Inc. Sizes of Ions Anions are larger than their parent atoms. Electrons are added and repulsions between electrons are increased.

28. © 2009, Prentice-Hall, Inc. Sizes of Ions In an isoelectronic series, ions have the same number of electrons. Ionic size decreases with an increasing nuclear charge.

29. Shielding As more PELs are added to atoms, the inner layers of electrons shield the outer electrons from the nucleus. The effective nuclear charge (enc) on those outer electrons is less, and so the outer electrons are less tightly held.

30. Ionization Energy This is the second important periodic trend. If an electron is given enough energy (in the form of a photon) to overcome the effective nuclear charge holding the electron in the cloud, it can leave the atom completely. The atom has been “ionized” or charged. The number of protons and electrons is no longer equal.

31. Ionization Energy The energy required to remove an electron from a gaseous atom is ionization energy. (measured in kilojoules, kJ) The larger the atom is, the easier its electrons are to remove. Ionization energy and atomic radius are inversely proportional. Ionization energy is always endothermic, that is energy is added to the atom to remove the electron.

32. Ionization Energy Energy required to remove an electron from a gaseous atom or ion. X(g) → X + (g) + e – Mg → Mg + + e – I 1 = 735 kJ/mol(1 st IE) Mg + → Mg 2+ + e – I 2 = 1445 kJ/mol(2 nd IE) Mg 2+ → Mg 3+ + e – I 3 = 7730 kJ/mol*(3 rd IE) *Core electrons are bound much more tightly thanvalence electrons.

33. Ionization Energy In general, as we go down a group from top to bottom, the first ionization energy decreases. Why? The electrons being removed are, on average, farther from the nucleus.

34. Ionization Energy In general, as we go across a period from left to right, the first ionization energy increases. Why? Electrons added in the same principal quantum level do not completely shield the increasing nuclear charge caused by the added protons. Electrons in the same principal quantum level are generally more strongly bound from left to right on the periodic table.

35. Ionization Energy

36. The Values of First Ionization Energy for the Elements in the First Six Periods

37. © 2009, Prentice-Hall, Inc. Trends in First Ionization Energies However, there are two apparent discontinuities in this trend.

38. © 2009, Prentice-Hall, Inc. Trends in First Ionization Energies The first occurs between Groups 2 & 3 In this case the electron is removed from a p- orbital rather than an s- orbital. The electron removed is farther from nucleus. There is also a small amount of repulsion by the s electrons.

39. © 2009, Prentice-Hall, Inc. Trends in First Ionization Energies The second occurs between Groups 15 and 16 The electron removed comes from doubly occupied orbital. Repulsion from the other electron in the orbital aids in its removal.

40. Which atom would require more energy to remove an electron? Why? Na Cl CONCEPT CHECK!

41. Which atom would require more energy to remove an electron? Why? Li Cs CONCEPT CHECK!

42. Which has the larger second ionization energy? Why? Lithium or Beryllium CONCEPT CHECK!

43. Successive Ionization Energies (KJ per Mole) for the Elements in Period 3

44. PES Photoelctron Spectroscopy Presented By, Mark Langella, APSI Chemistry 2013, PWISTA.com Method that provides information on all the occupied energy levels of an atom We can compare the energy required by a photon to knock an electron out of an atom (ionization energy), to the signal output to deduce how many electrons are at each energy level.

45. Methodology Very high energy photons, such as very-short- wavelength ultraviolet radiation, or even x-rays, are used in this experiment. The gas phase atoms are irradiated with photons of a particular energy. If the energy of the photon is greater than the energy necessary to remove an electron from the atom, an electron is ejected with the excess energy appearing as kinetic energy, ½ mv 2, where v is the velocity of the ejected electron. The speed of the ejected electron depends on how much excess energy it has received. Presented By, Mark Langella, APSI Chemistry 2013, PWISTA.com

46. Ionization Energy IE = the ionization energy of the electron KE = the kinetic energy with which it leaves the atom E photon = IE + KE or IE = E photon – KE kinetic energy of the electrons is measured in a photoelectron spectrometer photons of sufficient energy are used, an electron may be ejected from any of the energy levels of an atom every electron in each atom has an (approximately) equal chance of being ejected Presented By, Mark Langella, APSI Chemistry 2013, PWISTA.com

47. Spectrum for a large group of identical atoms, the electrons ejected will come from all possible energy levels of the atom because the photons used all have the same energy, electrons ejected from a given energy level will all have the same energy. Only a few different energies of ejected electrons will be obtained, corresponding to the number of energy levels in the atom. Presented By, Mark Langella, APSI Chemistry 2013, PWISTA.com

48. Photoelectron spectrum The photoelectron spectrum is a plot of the number of ejected electrons (along the vertical axis) vs. the corresponding ionization energy for the ejected electrons (along the horizontal axis) The “X” axis is labeled high to low energies so that you think about the XY intersect as being the nucleus. Presented By, Mark Langella, APSI Chemistry 2013, PWISTA.com http://www.chem.arizona.edu/chemt/Flash/photoelectron. html

49. Presented By, Mark Langella, APSI Chemistry 2013, PWISTA.com

50. Interpretations from the data: 1. There are no values on the y axis in the tables above. Using the Periodic Table and Table 1, put numbers on the y axis. 2. Label each peak on the graphs above with s, p, d, or f to indicate the suborbital they represent.. 3. What is the total number of electrons in a neutral potassium atom?

51. PES Question If a certain element being studied by a PES displays an emission spectrum with 5 distinct kinetic energies. What are all the possible elements that could produce this spectrum? Presented By, Mark Langella, APSI Chemistry 2013, PWISTA.com

52. Electron Affinity What does the word ‘affinity’ mean? Electron affinity is the energy change that occurs when an atom gains an electron (also measured in kJ). Where ionization energy is always endothermic, electron affinity is usually exothermic, but not always.

53. Electron Affinity Electron affinity is exothermic if there is an empty or partially empty orbital for an electron to occupy. If there are no empty spaces, a new orbital or PEL must be created, making the process endothermic. This is true for the alkaline earth metals and the noble gases.

54. Electron Affinity ++

55. Electronegativity Electronegativity is a measure of an atom’s attraction for another atom’s electrons. It is an arbitrary scale that ranges from 0 to 4. The units of electronegativity are Paulings. Generally, metals are electron givers and have low electronegativities. Nonmetals are are electron takers and have high electronegativities. What about the noble gases?

56. Electronegativity 0

57. Overall Reactivity This ties all the previous trends together in one package. However, we must treat metals and nonmetals separately. The most reactive metals are the largest since they are the best electron givers. The most reactive nonmetals are the smallest ones, the best electron takers.

58. Overall Reactivity 0