Chapter 5. The Periodic Table

5.1 Classification of the Elements Chapter 5: The Periodic Table 5.2 The Periodic Law Dimitri Mendeleev (1869) arranged elements in a table in order of increasing mass. He put elements with similar properties in the same column. The Periodic Law H. G. J. Moseley arranged elements in order of increasing atomic number. Neils Bohr’s introduction of electron energy levels altered the shape of the Periodic Table. Periodic Law – elements in the same column have similar properties The modern Periodic Table shows s, p, d and f sublevels. 5.3 Organization of the Periodic Table A horizontal row is called a period. A vertical column is called a group. Elements in the same group exhibit similar properties. Main Group (Representative) Elements: A Group Transition Metals: B Group 1 Group Names of Elements Group Names of Elements Group IA: alkali metals  Transition Metals (B group) Group IIA: alkaline earth metals  Inner Transition Metals (beneath the main body Group VIIA: halogens Group VIIIA: noble gases of Periodic Table) • lanthanide series: Ce-Lu, also called rare earth metals • actinide series: Th-Lr, all radioactive elements Note: all elements > 93 are man-made in particle accelerators. Test Yourself The Alkali metal in the third period is ____ The Halogen in the 2nd period is _____ The Noble gas in the fourth period is ____ 5.4 Atomic Size  Atomic Size: average distance from nucleus to outermost electron  Increases down a group/column  As we go , there are more energy levels and radium gets bigger  Increases from right to left () across a period/row Atomic Size  As we go , there are more protons and more positive nucleus. As positive charge increases, the electrons are pulled closer, so atoms get smaller. This is called the effective nuclear charge.  Trend is like a snowman that fell to the right  Trend increases  and  2 Atomic Size Which is largest? A) Cl, Se, or Te B) Mg, Si, or S Metallic Character 5.5 Properties of Elements  Increases from right to left   metals are on left-hand side  Elements in the same group typically have  Increases down a group   For groups IVA and VA, we go from  For example, nonmetals (C & N) at the top to metals (Sn & Bi) at the bottom. similar chemical behavior and physical properties  Na reacts with Cl2 to form NaCl  K reacts with Cl2 to form KCl  Same snowman trend!  Trend increases  and  Properties of Elements Ex. 1: Sodium reacts violently with water. Which other element(s) also will? K Mg Al Li P 5.7 Valence Electrons  Core electrons: inner e- in filled electron levels  Valence electrons: s and p e- in the outer unfilled level  Valence electrons form chemical bonds and Ex. 2: Chlorine reacts with hydrogen to form HCl. Give the formula when fluorine reacts with hydrogen. _____ dictate and element’s properties and chemical behavior. 3 Valence Electrons 5.8 Writing Electron Dot Formulas  For Group A elements: Group # = # of valence electrons IA = 1 v.e. IIA = 2 v.e. IVA = 4 v.e. VA = 5 v.e. VIIA = 7 v.e. VIIIA = 8 v.e. IIIA = 3 v.e. VIA = 6 v.e.  How many valence electrons do the following • elements have? Si? Kr? P? Na? Ca? S? 1. Write down the element symbol. 2. How many valence electrons? 3. Assume the atom has four sides, and put electrons with one electron per side before pairing. There are a maximum of 2 e– on each side! Let’s try some on the board! 5.9 Ionization Energy  Ionization energy: energy required to remove an electron from a neutral atom.  IE increases up a group   easier to remove electrons from a larger atom (e–’s are further from the nucleus)  IE increases left to right   Metals have low IE’s because they want to lose electrons to be like a noble gas  Nonmetal elements have high IE’s because they would rather gain electrons. These trends are the opposite of the snowman Ionization Energy 5.10 Ion Formation Which element in each group has the largest ionization energy?  Atoms gain or lose electrons to have same A) Cl, P, or Ca?  Noble gas: charge = 0; very stable, B) Li, Rb, or Be? number of valence electrons as a noble gas. chemically inert, happy with 8 valence electrons (full s and p)  Except He has full 1s only = 2 e- (He does not have 8 valence electrons!) 4 Ion Formation Ion Formation  Metals lose electrons and become  Nonmetals gain electrons and become positively charged ions called cations Group Charge Example Group IA metals +1 Group IIA metals +2 Group IIIA metals +3 Li+ Mg2+ Al3+ negatively charged ions called anions Group Group VA Group VIA Group VIIA nonmetals nonmetals nonmetals Charge Example –3 N3– –2 O2– –1 F– • magnesium ion can be shown as Mg+2 or Mg2+. • The 1 is omitted for +1 or –1 What charge will ions of these elements have? lithium magnesium sulfur bromine nitrogen aluminum Electron Configuration of Ions  + charge means less e- than the atom (cations lose electrons)  – charge means more e- than the atom (anions Isoelectronic Iso- = same -electronic = electrons Isoelectronic = 2 elements that have the same number of electrons and therefore the same electron configurations! Electron Configuration of Ions 1) Na+ is like Ne Na atom 1s22s22p63s1 Na+ (isoelectronic w/?) gain electrons)  Write the electron configuration of the neutral atom first, then remove the correct number of electrons for cations, or add electrons for anions. 2) Mg2+ is like ? Mg atom 1s22s22p63s2 Mg+2 (isoelectronic w/?) 3) F- is like ? F atom 1s22s22p5 F- (isoelectronic w/?) 5 Ion Isoelectronic with Electron which Noble gas? Configuration Answers  They are all the same K+  1s22s22p63s23p6 P3-  All isoelectronic with Argon ClCa2+ Isoelectronic Ions Example 2: Circle all the ions below that are isoelectronic with neon: Coulomb’s Law  Law of physics published in 1783  Basically describes how opposites attract.  This is VERY important in chemistry as ions of opposite charge attract  The electrostatic attraction between two charges is proportional to the charge magnitude (q) and inversely proportional to the distance (r) squared.  F = ke(q1q2 / r2) Ch. 5 Self Test, p. 159  Try numbers 2-4, 7-13, 15-17  Answers in Appendix J 6