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The Periodic Properties of the Elements

The Periodic Properties of the Elements

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The Periodic Properties of the Elements

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  1. The Periodic Properties of the Elements By Lauren Querido, Chris Via, Maggie Dang, Jae Lee

  2. The Founders of the Periodic Table Luthar Meyer Dmitri Mendeleev http://nuclphys.sinp.msu.ru/persons/images/mendeleev.gif http://chemheritage.org/classroom/chemach/images/lgfotos/04periodic/meyer-mendeleev2.jpg

  3. 7.1 Developing the Periodic Table • Dmitri Mendeleev (1869)- and Luthar Meyer Published very similar documents to classify the elements. And were the first to make the modern periodic table • Used chemical and physical properties to classify • Henry Moseley (1887-1915)- Developed concept of atomic numbers • Found that frequency increases as the atomic mass increases

  4. 7.2 Electron Shells and Size of Atoms • Electron Shells in Atoms • Gilbert N. Lewis – electrons are arranged in shells surrounding the nucleus. • Atomic sizes-http://grandinetti.org/Teaching/Chem121/Lectures/PeriodicTrends/assets/radiitable.gif

  5. Bonding Atomic Radius- the distance between the center of two bonding atoms http://www.chembook.co.uk/fig13-1.jpg

  6. Practice Problem #1 • Predict the lengths of C-S, C-H, and S-H bonds in this molecule • Radius of C = 0.77 Å • Radius of S = 1.02 Å • Radius of H = 0.37 Å • When determining the bonding radius, you add the radius of the bonding atoms together

  7. Answer to Practice Problem #1 • C-S bond length = radius of C + radius of = 0.77 Å + 1.02 Å = 1.79 Å • C-H bond length = 0.77 Å + 0.37 Å = 1.14 Å • S-H bond length = 1.02 Å +0.37 Å = 1.39 Å

  8. 7.2 continued • When moving across a row, the number of core electrons stay the same but the nuclear charge increases • The effective nuclear charge increases even though the quantum number remains the same • Shielding is the process of blocking the protons effective charge on the outermost electrons

  9. 7.3 Ionization Energy • Ionization Energy – to remove an electron from the ground state • Second Ionization – removing the 2nd electron from the ground state • I1<I2<I3 and so forth; It increases in magnitude • The greater effective nuclear charge, the greater the ionization energy

  10. 7.3 cont.. • There is a sharp increase in ionization energy when an inner shell electron is removed • Periodic Trends • Within each row, the ionization energy increases with atomic number • Within a group, the ionization energy generally decreases with increasing atomic number

  11. 7.3 cont.. 3. The ionization energy of transition elements & f-block metals increase slowly as you read from left to right. • The transition in ionization energy are affected by how strong an electron is attracted to an atom • It is affected by the effective nuclear charge and the average distance from the nucleus.

  12. 7.3 • The irregularities are explained through the periodic table • Electrons in the s orbital are more effective at shielding than in the p orbital

  13. 7.4 Electron Affinities • Positive ionization energy = energy put into atom in order to remove electrons • Electron affinity = attraction of change in energy when the electron is added • Most atoms = energy is released when electron is added • A positive electron affinity, an ion will not form

  14. 7.4 cont.. • On the periodic table, electron affinity becomes negative towards halogen (closest to being stable) • The electron affinity does not change when they move down a group (noble gases)

  15. http://www.meta-synthesis.com/webbook/35_pt/best_PT-form.jpg

  16. Element Classificationhttp://www.elementsdatabase.com/Images/periodic_table1.gif

  17. 7.5 Metals, Nonmetals, and Metalloids • Metals • Tend to have low ionization energies and lose electrons when they undergo chemical reaction • Most metal oxides are basic oxides that dissolve in water react to for metal hydroxides Metal Oxide + Water  Metal Hydroxide • Metal oxides show their basicity by reacting with acids to form water and salts Metal Oxide + Acid  Salt + Water

  18. 7.5 • Characteristics of Metals • Have a shiny luster • Various colors • Solids are malleable and ductile • Good conductors of heat and electricity • Most metal oxides are ionic solids that are basic

  19. 7.5 • Nonmetals • Tend to gain electrons and become anions Metal + Nonmetal  Salt • Most nonmetal oxides are acidic oxides that dissolve in water react to form acids Nonmetal Oxide + Water  Acid • The acidity of nonmetal oxides is shown by the fact they dissolve in basic solutions to form salts Nonmetal Oxide+ Base  Salt + Water

  20. 7.5 • Characteristics of Nonmetals • Do not have a luster • Various colors • Solids are usually brittle; some are hard, and some are soft • Poor conductors of heat and electricity • Most nonmetallic oxides are molecular substances that form acidic solutions

  21. 7.5 • Metalloids • Have properties intermediate between nonmetals and metals http://www.rkm.com.au/METALLOIDS/metalloid-images/METALLOID-SILICON-500.jpg

  22. 7.6 Group Trends for the Active Metals • Group 1A: The Alkali Metals (most active) • Metallic Characteristics • Silvery • Metallic luster & high thermal • Electrical conductivities • Have low densities & melting points • Very reactive b/c they want to lose 1 electron to form ions with a 1+ charge so it becomes more stable

  23. 7.6 cont.. • As you move down a group • Atomic radius increases • 1st ionization energy decreases

  24. 7.6 cont.. • Group 2A: The Alkaline Earth Metals • Properties of Alkaline Earth Metals • Harder • More Dense • Melt at higher temps • Highly Reactive • Compared to alkali Metals, Alkaline Earth metals.. • Have lower 1st ionization energies • Are less reactive

  25. 7.7 • Group 6A • Oxygen is a colorless gas at room temperature while all the other elements in this group are solid. • Oxygen has two main forms: 02=“oxygen” and 03=“ozone”. • This is an example of an allotrope, it has different forms of the same element.

  26. 7.7 • The most stable form of sulfur is S8, It is a yellow solid. • All of the elements in this group have the tendency to gain electrons form other elements. • http://www.science.uwaterloo.ca/~cchieh/cact/fig/s8.gif

  27. 7.7 • Group 7A: Halogens • Halogens is named Greek words, “halos” and “gennao” meaning salt formers. • Fluorine and Chlorine are gases, bromine is a liquid, and iodine is a solid at room temperature. • These elements melting and boiling points increase with atomic number. • These elements have highly negative electron affinities because they have the need to gain electrons from other elements.

  28. The Fluorine atom is very reactive! http://www.chemistryland.com/ElementarySchool/BuildingBlocks/FluorineAttracts.jpg

  29. 7.7 • Group 8A: Noble Gases • All of the elements are nonmetals at room temperature and they are monatomic • They are very unreactive because they have completely filled s and p orbitals. • They also have very large 1st ionization energies.

  30. That’s All Folks