AP Chemistry Review of The Periodic Table

1. AP Chemistry Review of The Periodic Table What is a Trend? A general direction in which something tends to move. A general tendency or inclination. What does Periodic mean? Recurring or reappearing at regular intervals. What are some examples of natural, periodic trends? Moon phases, ocean tides, water temperature in the northeast…. How does the Periodic Table of Elements illustrate Periodic Trends in Physical and Chemical Properties? Use Reference Table S to examine periodic trends in 1st Ionization Energy and Atomic Radius when compared with atomic number. What is 1st Ionization Energy and what does it tell us about an atom? 1st Ionization Energy is the Energy required to remove one mole of electrons from one mole of atoms (Thus, each atom becomes an ion). It tells how strongly an atom attracts its outermost electrons.

2. State the general trend and describe the periodic character. Generally, as atomic no. increases 1st Ionization Energy also increases, but periodically this trend “starts over again”. How does the Periodic Table show this periodic character? The atoms that start and end the trend also start and end the Rows called Periods.

3. State the general trend and describe the periodic character. Generally, as atomic no. increases, the atoms become smaller!? Periodically, this trend also “starts over again.” How does the Periodic Table show this periodic character? Once again, the atoms that start and end the trend also start and end the Rows!

4. Illustrate the trends in 1st Ionization Energy and Atomic Radius on the blank Periodic Table below. Strongest Hold on Outer Electrons 1st Ionization Energy Increases 1st Ionization Energy Decreases 1st Ionization Energy Decreases Atomic Radius Decreases Atomic Radius Increases Atomic Radius Increases Weakest Hold on Outer Electrons Consult Ref. Table S and the two previous graphs and Illustrate the same trends going Down the Columns at the two ends of the Table. Indicate the location of the atoms with the strongest hold and the atoms with the weakest hold on their outer electrons. How do you know? Highest 1st Ionization Energy = Strongest hold and visa versa

5. Why do the atoms have a stronger hold on outer electrons as one moves Across a Period (row)? The number of Protons or Nuclear Charge increases while the outer electrons are in the same shell or energy level (Bohr model). Why do atoms have a weaker hold on outer electrons as one moves Down a Column? Although the nuclear charge increases, the outer electrons are one shell farther from the nucleus. Force is proportional to 1/d2 (inverse squared). As distance increases force decreases a lot!! As one moves across a period, the atoms have more protons, more neutrons and more electrons. Why, then, do they become smaller??! Nuclear Charge increases while the outer electron shell (distance) remains the same. The outer electron feels a stronger pull. If having more protons in the nucleus makes atoms smaller, why then, do they become larger as one moves Down a Column? As one moves down the column from one row to another, the distance to the outer electron shell increases.

6. A Brief History of the Development of the first Periodic Table Additional Web Sites: Interactive Periodic Table. Very Good http://www.chemsoc.org/viselements/pages/pertable_fla.htm  WebElements Interactive Periodic Table. Good http://www.webelements.com/webelements/elements/text/periodic-table/phys.html  CE&N Interactive Periodic Table http://pubs.acs.org/cen/80th/elements.html History of the Development of the Periodic Table. Very Brief. Starts with Dimitri Mendeleev in 1869 http://www.upei.ca/~physics/p221/pro00/periodicTble/page2.html

7. Metals vs. Nonmetals Many are gases or liquids at room temp. Most are solids at room temp. Good Conductors of Heat and Electricity Poor Conductors of Heat and Electricity Solids have Luster or Sheen Solids tend to be Dull Malleable: can be Hammered into thin Sheets Solids tend to be Brittle and will shatter if hit with a hammer Ductile: can be Drawn into thin Wires Metalloids show a blend of characteristics

8. Periodic Trends in Metallic vs. Nonmetallic Character What is the general trend in metallic vs. nonmetallic character across a period? Across a row, elements change from more metallic to more nonmetallic. How does the presence of the Metalloid group (B, Si, Ge, As, Sb, Te) illustrate this general tend? The change in character is a gradual change from metallic to nonmetallic. What do the stairs indicate? The stairs indicate a general boundary between metals and nonmetals. What element’s location seems exceptional? Hydrogen is a nonmetal.

9. Atomic Electron Configuration and the Periodic Table The Bohr Electron Configuration indicates the number of electrons in each shell or principle energy level of the atom What Similarities in electron configuration exist within Chemical Families (columns) in the Periodic Table? Generally, elements in the same Family have the same number of outer electrons. What general Differences in electron configuration exist between Metals and Nonmetals? Metals tend to have only a few outer electrons while nonmetals tend to have nearly full outer energy levels. The Noble Gases (Group 18) are the most chemically stable or unreactive elements. What might account for their stability? Noble Gas atoms all have filled outer shells with 8 electrons (except Helium).

10. Atoms and the Stable Ions They Tend to Form The Selected Oxidation States “give a clue” as to what Stable Ion an atom may form (More in Unit 8). There is a natural tendency for all things to go to the most stable (and lowest energy) state. The Noble Gases are chemically the most stable elements. What do metals and nonmetals tend to do to become more stable? Why? Metal atoms tend to lose their outer electrons and form positive ions while nonmetal atoms tend to gain more outer electrons and form negative ions in order to fill their outer electron shells like the nearest Noble Gas. Cation (“cat-eye-on”) is a positive ion. Think ca+ion. Anion (“an-ion”) is a negative ion. Think anion = anegative ion. How does the size of an atom compare with the cation or anion it forms? Why? Cations are smaller and Anions are larger. Because 

11. Atomic Orbitals have Specific, Yet Cloudy Shapes

12. Orbitals Represent Sublevels Within Principle Energy Levels 3d after 4s??????

13. Electrons Fill the Orbitals in a Specific Order

14. Effective Nuclear Charge and Periodic Properties The Effective Nuclear Charge, Z*, indicates “What the outermost electron feels.” How strongly the outermost electron feels the attraction to the nucleus depends on: (1) Actual Nuclear Charge = the number of protons (2) Shielding(repulsion) by (a) inner electrons and (b) electrons in the same shell (“crowding”) (3) Distance from the nucleus determined by the energy level. Many physical and chemical characteristics of atoms and ions can be explained by understanding the effective nuclear charge. See notes from class discussions Compare each of the atoms in the second period: Compare atoms of the same family: Compare atoms with their cations: Metal cations are one shell smaller. Nonmetal anions have more “crowding”. Compare atoms with their anions:

15. Comparative Sizes of Atoms and Ions When comparing Atoms to Atoms look up the Atomic Radius in Ref. Table S. Be sure of whether the question asks for largest or smallest. When comparing Atoms to Ions or comparing Ions to Ions (1) Draw the Bohr Electron Configuration Model and (2) Determine the Effective Nuclear Charge. Which particle is largest? C vs. N Ge vs. Se Mg ion vs. Al ion K ion vs. Cl ion O2- vs. F- K+ vs. Ca2+ Which particle is smallest? S vs. Br Ne vs. Na+ Br - vs. Kr Li+ vs. Be2+ Na ion vs. F ion N Br Se Na+ Mg2+ Kr Cl- Be2+ O2- Na+ K+

16. Names of Common Chemical Families Indicate the column for each of the following:Alkali Metals, Alkaline Earth Metals, Transition Metals, Halogens, Noble Gases

17. Valence Electrons and Lewis Electron Dot Structures Valence Electrons are the electrons in the outermost energy level or outermost electron shell. From Valiant or Strong as these electrons are involved in the bonds that hold atoms together. Lewis Electron Dot Structures indicate an atom’s valence electrons. The Chemical Symbol indicates the nucleus and all inner electrons. One dot is placed around the chemical symbol for each valence electron. There are four places to put the dots. Each place can hold up to two dots. Give the Lewis Electron Dot Structure for each of the first 20 Elements.

20. Metals and Nonmetals Form Ionic Bonds What do Metals tend to do? Metals tend to lose their valence electrons and form stable, positive ions. What do Nonmetals tend to do? Nonmetals tend to gain enough electrons to fill their valence shell and form stable, negative ions.

21. When Metals and Nonmetals are allowed to react, (1) Valence electrons are Transferred from the metal to the nonmetal. (2) The atoms then become Ions with opposite charges. (3) Their Electrostatic Attraction form Ionic Bonds. The ions arrange themselves in a neat, orderly, repeating pattern (crystal) with a specific ratio.

22. Ionic Compounds form Ionic Crystals The Stable Geometric Structure gives Ionic Compounds High Melting and Boiling Points. Most Ionic Compounds are soluble in water ….. At least to some degree. Ionic Compounds are Electrical Conductors only when Melted or when Dissolved in Water. The ions must be free to move for conduction.

23. Ionic Compounds and Chemical Formula Ionic Compounds are Neutral. Therefore, what can you infer about the ratio of ions in an Ionic Compound? There must be enough positive and negative ions that the charges “cancel out” The Chemical Formula of ionic compounds indicates the simplest ratio of ions in the compound, positive ion first, negative second. Give the correct Chemical Formula for each of the following: Ex. Calcium Chloride: CaCl2 Potassium Sulfide: K2S Magnesium Oxide: MgO Aluminum Iodide: AlI3 Aluminum Oxide: Al2O3 The “-ide Rule”: For monatomic anions, change the ending to –ide. Polyatomic Ions: See Ref. Table E. Use parenthesis if more than one is needed in the formula Sodium Carbonate Calcium Nitrate Ammonium Phosphate Na2CO3 Ca(NO3)2 (NH4)3PO4

24. For atoms with more than one possible charge, the Stock System uses a Roman numeral to indicate the charge on the ion. Write the Chemical Formula of the following: Copper(II) Nitrate Iron(III) Chloride Lead(IV) Sulfite Cu(NO3)2 FeCl3 Pb(SO3)2 One can determine the charge (and possibly the group) of an unknown ion from the ionic formula. X2O MgX2 XS BeX Na3X X(OH)3 Fe2O3 Fe is….? X1+ X-1 X2+ X2- Fe3+ X3- X3+ For each example above, identify the Group that element X belongs in and Name the Group if it has a “common” name.