Periodic Properties of the Elements

1. Periodic Properties of the Elements Suh Kwon Abegim Undie _ Emily Scheerer Justin Green

2. 7.1 Development of the Periodic Table Elements have been being discovered since the beginning of time. As the number increased, scientists needed a way to organize the elements In 1869 Dmitri Mendeleev and Lothar Meyer both published classifications that noted the periodic similarities between elements.

3. 7.1 Development of the Periodic Table At the time, there was no knowledge of atomic numbers However, both scientists arranged elements with increasing weight. Medeleev is given the credit because he pronounced his ideas more. For example, he predicted the existence and properties of both gallium and germanium. http://www.docbrown.info/page03/3_34ptable/PTmendeleev1869.gif

4. 7.1 Development of the Periodic Table In 1913, Henry Mosley developed the idea of atomic numbers. Determined that each element produced unique X-ray frequencies, assigning each element a number based on its X-ray frequency. He then identified that the atomic number was equal to the number of protons and electrons in the atom. This clarified problems with the weight arrangement, allowing them to find ‘holes’. Can you find an example? (Ar and K)

5. 7.2 Electron Shells & Sizes of Atoms What shape does an atom have??? According to the quantum mechanical number, an atom does not have a defined shape. http://physics.uwstout.edu/geo/bedtime/graphics/atom.jpg

6. 7.2 Electron Shells & Sizes of Atoms As you move down the periodic table, n changes. Gilbert Lewis: suggested that electrons in atoms are arranged in spherical shells. Radial electron density: the probability of finding the electron at a particular distance from the nucleus.

7. 7.2 Electron Shells & Sizes of Atoms At certain distances, RED shows maxima. This indicates higher probability of finding electrons. How many maximums: due to electrons that have the same n value. Example: Helium has one maximum because it has one n value (1) Argons has three maximums because of it’s three n values. (1,2,3)

8. http://upload.wikimedia.org/wikipedia/commons/e/e5/Periodic_table_of_elements_showing_electron_shells.pnghttp://upload.wikimedia.org/wikipedia/commons/e/e5/Periodic_table_of_elements_showing_electron_shells.png

9. 7.2 Electron Shells & Sizes of Atoms Defining atomic size: Either nonbonding atomic radius(Van der Waals radii) When 2 atoms collide and bounce off of each other. It’s then the distance from the nucleus to the outer edge OR bonding atomic radius (covalent radii) When two atoms collide and an attractive interaction leads to a chemical bond. The distance between the two nuclei, shorter than the non-bonding atomic radius. Why?

10. 7.2 Electron Shells & Sizes of Atoms Scientists have developed a variety of methods for measuring these distances. From observations of these methods, each element is assigned a bonding atomic radius. Example: Iodine According to observation, the distance separating the nuclei from the electrons is 2.66 angstroms & therefore the bonding atomic radius is 1.33 angstroms. (cut in half to find radius of one)

11. 7.2 Electron Shells & Sizes of Atoms To find atomic radii of a bond Add atomic radii of the two bonding atoms (NOTE: This is NOT 100% accurate) ~ See Example 7.1 on page 232.

12. 7.2 Electron Shells & Sizes of Atoms • Periodic Trends of Atomic Radii: • Increases from top to bottom • Increasing n value means a larger orbital, which means a longer radius. • Decreases from left to right • Zeff increases, pulling electrons closer, making the radius smaller See Example 7.2 on page 233 http://grandinetti.org/Teaching/Chem121/Lectures/PeriodicTrends/index.html

13. 7.2 Electron Shells & Sizes of Atoms http://grandinetti.org/Teaching/Chem121/Lectures/PeriodicTrends/index.html http://chemistintheory.blogspot.com/2008_02_01_archive.html

14. 7.3- Ionization Energy • Energy required to remove an electron from the ground state of the atom or ion in standard conditions http://www.avon-chemistry.com/p_table_sus_ion.jpg

15. 7.3- Ionization Energy • The greater the Ionization energy, the more difficult it is to remove an electron • As the ionization energy increases, the atomic number goes up • I1  first ionization • Energy needed to remove the first electron from a neutral atom • I2 second ionization • Energy needed to remove 2nd electron • And so forth, for successive removals of additional electrons

16. 7.3 – Ionization Energy • Periodic Trends • I1 increases with increasing atomic number • Alkali metals show lowest ionization energy in each row • Noble gases are the highest ionization energy • Within each group (up and down), ionization energy decreases with increasing number • HE > Ne > Ar > Kr > Xe • Representative elements (sublevels “s” and “p”) show a larger range of values of I1 • F-blocks only show a small variation of values of I1

17. 7.3 – Ionization Energy • Factors that affect how strongly an electron is attracted to an atom • Nuclear charge (higher charge means there are more protons) • Average distance of the electron from the nucleus

18. 7.3 – Ionization Energy • For instance, attraction increases when nuclear charge increases and distance decreases • As the nuclear charge increases, there are more and more protons • The protons have a strong attraction for the electrons, so the distance decreases • Thus, the ionization energy increases

19. 7.3 – Ionization Energy • Ionization energy measures the energy changes associated with removing electrons from an atom • Positive value of ionization energy tells us that energy must be put into an atom to remove the electron

20. 7.4 – Electron Affinities • Measures the attraction of the atom for the added electron • The more negative the electron affinity, the greater the attraction between the atom and electron • An electron affinity that is > 0 shows us that the negative ion is higher in energy than the separated atom and electron

21. 7.4 – Electron Affinities • Difference between Ionization and Electron Affinities • Ionization  measures the ease with which an atom loses an electron • Electron affinities  measures the ease with which an atom gains an electron

22. 7.4 – Electron Affinities • Periodic Trends • Halogens have the most-negative electron affinities • The addition of an electron to a noble gas requires that the electron reside in a higher-energy sub shell • Occupying a higher-energy sub shell (1s, 2s, 2p, 3s…) is unfavorable, so the electron affinity is positive

23. 7.5 - Metals, Nonmetals, and Metalloids • Properties of atoms include atomic radii, ionization energies, and electron affinities • No elements exist in nature as an individual atom except for the Noble Gasses.

24. Periodic Table Grouping • The periodic table is grouped into meals, nonmetals, and metalloids • Metals take up the top left and middle portions of the periodic table, nonmetals occupy the right side (and hydrogen), and metalloids are located between the two http://www.schenectady.k12.ny.us/users/title3/Future%20Grant%20Projects/Projects/periodictable/Regions.jpg

25. Metallic Character • The more an element exhibits the physical and chemical properties of a metal, the greater it’s metallic character • Metallic character increases as you go down the periodic table and increases as you go from right to left http://www.global-b2b-network.com/direct/dbimage/50357240/Electrolytic_Manganese_Metal.jpg

26. Metals • Characteristics • Shiny • Conduct heat and electricity • Malleable • Ductile • Solid at Room Temperature (except mercury) • Higher than room temperature melting points • Low ionization energy and form positive Ions easily • Oxidize as they undergo reactions http://www.ndt-ed.org/EducationResources/CommunityCollege/Materials/Graphics/MixedMetals(mayFranInt.).jpe

27. Metals Cont. • Alkali metal ions always have a charge of +1 • Alkaline earth metal ions always have a charge of +2 • Transition metal ions are mostly +2, but they do have +3 and +1 ions • Some transition metals have multiple charges due to their position on the periodic table • Compounds between metals and nonmetals tend to be ionic • 2Ni(s) + O2(g) 2NiO(s) • Most metal oxides are basic • Metal oxide + water metal hydroxide • Na2O(s) + H2O(l) Ca(OH)2(aq) • Metal oxide + acid salt + water • NiO(s) + 2HCl(aq) NiCl2(aq) + H2O(l)

28. Nonmetals • Vary in appearance • Poor conductors of heat and electricity • Melting points are generally lower than metals • Have seven diatomic molecules (Br2,I2,N2,Cl2,H2,O2,F2) • React with metals to form salts • Metal + Nonmetal Salt • 2Al(s) + 3Br2(l) 2AlBr3(s) • Most nonmetal oxides are acidic oxides • Nonmetal oxide + water acid • Dissolve in basic solutions to form salts • Nonmetal oxide + base salt + water http://www.indiamart.com/sujaychemicals/pcat-gifs/products-small/liquid-20bromine_10438902.jpg

29. Metalloids • Have properties intermediate between those of metals and nonmetals. They may have some characteristics of metallic properties but lack others • Semiconductors • Silicon http://1366tech.com/v1/images/stories/site/silicon.jpg

30. Group Trends for the Active Metals • Elements in a group posses general similarities • Trends occur when you move through a group or from one group to another

31. Group 1A: The Alkali Metals • Soft metallic solids • low melting points and low densities • Atomic radius increases as you travel down the column • Ionization energy decreases as you travel down the column • Very reactive • Combine directly with nonmetals • react vigorously with water • 2M(s) + H2O(l) 2MOH(aq) + H2(g) • Reacts with oxygen to produce a metal oxide • when placed in a flame they create different colors http://content.tutorvista.com/chemistry_11/content/us/class11chemistry/chapter12/images/img19.gif

32. Group 2A: The Alkaline Earth Metals • Harder, more dense, and have higher melting points than the elements of the 1A column • Less reactive than the Alkali metals • Calcium and elements below it will readily react with water at room temperature whereas magnesium will only react with steam and Beryllium will not react at all with water • Because of their relatively high reactivity, the alkaline earth elements are invariably found in nature as compounds of the 2+ ions http://www.learner.org/interactives/periodic/images/alkaliearthmetals.gif

33. Group Trends for Selected Nonmetals

34. Hydrogen http://library.thinkquest.org/C005858/hydrogen2.jpg

35. Hydrogen • First element in the periodic table • Does not truly belong to any family • Occurs as a colorless diatomic gas, H2 under most conditions • Owing to complete absence of nuclear shielding, ionization energy of hydrogen is very high (1312 kJ/mol)

36. Hydrogen (Cont) • Generally reacts with other nonmetals • Reactions can be exothermic • Hydrogen reacts with active metals to form solid metal hydrides

37. OXYGEN http://bp1.blogger.com/_83R26vRzoNM/Ruo5uekM8zI/AAAAAAAAAWA/___YOi9YZt4/s320/oxygen1.jpg

38. Group 6A: Oxygen Group • Metallic character increases as you go down • Oxygen is a colorless gas at room temperature, all others are solid • Oxygen is encountered in two molecular forms, O2 and O3 (O2 is most common) • O3 form is called ozone • Two forms of oxygen called allotropes

39. Oxygen Group (Cont) • Oxygen has a tendency to attract electrons from other elements (called oxidization) • Formation of nonmetal oxides is very exothermic and energetically favorable • Usually creates the stable oxide, the O2- ion

40. Oxygen Group (cont) • Second most important member of group 6A is sulfur • Sulfur has a tendency to gain electrons from other elements to form sulfides

41. Halogen http://www.lightbulbs2u.com/images/halogen_group.jpg

42. Group 7A: The Halogens • Halogens comes from Greek words halos and gennao, meaning “salt formers” • As we go from 6A to 7A, nonmetallic behavior of elements increases • Melting and boiling points increase with increasing atomic number

43. Halogens (cont) • Halogens have highly negative electron affinities • Halogens have a tendency to gain electrons from other elements to form halide ions • Halogens react directly with most metals to form ionic halides • Also react with hydrogen to form gaseous hydrogen halide compounds • These compounds are all very soluble

44. Group 8A: The Noble Gases • All nonmetals that are gases at room temperature • All monoatomic (consiste of single atoms rather than molecules)

45. Noble Gases (cont) • Completely filled s and p subshells • this very stable electron configuration makes them very unreactive • Only the heaviest noble gases form compounds, and only with very active nonmetals such as fluorine