CHAPTER 5 THE PERIODIC LAW What types of useful information can you find on the Periodic Table? I. History of the Periodic Table A. Before the Periodic Table was invented, about 63 elements were known. However, they were not organized and only random properties were known about each of the elements. Scientist (who are always looking for patterns) wanted to organize these. B. Dmitri Mendeleev he made cards for all 63 known elements and included some of their properties. He tacked them up on his laboratory wall and searched for patterns. Mendeleev arranged the elements by increasing atomic mass. When he did this, he found a periodic pattern of chemical properties. In a few cases, he had to put elements out of atomic mass order so their properties would line up with other elements of similar properties. Mendeleev s Questions: 1. Why are some elements out of atomic mass order? 2. What causes this periodicity of chemical properties? Mendeleev lost the Nobel Prize3 in the year before he died by one vote to the discoverer of fluorine. Although his table is now the most valuable tool in chemistry, he never won the Nobel Prize for it. C. Answer to Mendeleev s 1 st Q: In 1911, Henry Moseley (who worked with Rutherford) discovered that the positive charge of the nucleus increases by one unit as you move from one element to the next across Mendeleev s table. He decided that the amount of positive charge in the nucleus was based on the number of protons, and named this the atomic number. The Periodic Law - The physical and chemical properties of the elements are periodic functions of their atomic numbers. 1
D. The Modern Periodic Table - elements are arranged by increasing atomic number, so that elements with similar properties are in the same column. 1894 1900 discovery of the noble gases (group 18) Early 1900 s lanthanides (rare-earth elements) Period 6, #58 71) Early 1900 s actinides (al radioactive) Period 7 #90 103 The answer to Mendeleev s 2 nd Q: The reason for periodicity is the arrangement of electrons around the nucleus of each atom. Nuts and Bolts Lab Mendeleev s P.T. Puzzle Lab II. Electron Configuration and the Periodic Table Outer shell e - configurations can be used to explain the chemical properties of the elements. A. Period numbers and e - configs: Ex: Find the e - config. of: H 1s 1 outer level: n=1 He 1s 2 Li 1s 2 2s 1 Be 1s 2 2s 2 outer level: n=2 B 1s 2 2s 2 2p 1 Ne 1s 2 2s 2 2p 6 Pattern? **Period number (1-7) = the outer most occupied energy level in the ground state of that atom. So, what is the outermost occupied level in a K atom? 4 a Fr atom? 7 a P atom? 3 What period would an atom with the e - config. 1s 2 2s 2 2p 6 3s 1? 3 How about one with the config. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8? 4 2
B. group number and e - configuration 1. s-block elements Look at Li 1s 2 2s 1 Na 1s 2 2s 2 2p 6 3s 1 K 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 1 outer shell e - in an s orbital. Be 1s 2 2s 2 Mg 1s 2 2s 2 2p 6 3s 2 Ca 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 2 outer shell e - in an s orbital. Notice any pattern? ** s-block elements: last electron goes into an s orbital. Also, the group number is the same as the number of outer shell e -. 2. p-block elements (Groups 13 18) have their last e - going into a p orbital Ex: B 1s 2 2s 2 2p 1 3 outer shell e - Group 13 C 1s 2 2s 2 2p 2 4 outer shell e - Group 14 Ne 1s 2 2s 2 2p 6 8 outer shell e - Group 18 **Rule for s and p block elements (Groups 1-2, 13 18, called the main group elements The last digit of the group # tells the # of outer shell e -. So, how many outer shell e - does Al have? 3 Se Br 6 7 3
Draw the electron dot notation for S For Fr For Rn How many outer shell e - are in a Bi atom? 5 What energy level (n) are they in? 6 P-block elements vary in properties within the group, since the line of demarcation cuts through them. Groups 1: the alkali metals (Li, Na, K, Rb, Cs, Fr) --Highly reactive -- not found as free (uncombined) elements in nature; always found in compounds -- all react strongly with water and air, so they are stored under kerosene. -- all can be cut with a knife Hydrogen makes up 76% of universe. Placed above Group 1 (not init, even though it has 1 outer e - ) because a. It is not a metal b. It is chemically unlike the Group 1 elements Group 2: the alkaline-earth metals (Be, Mg, Ca, Sr, Ba, Ra) -- less reactive than Group1, but still founds only as compounds in nature. Helium - even though it has 2 outer shell e -, it is placed in Group 18 rather than Group 2, because it is nonreactive like the other noble gases. Group 17 the halogens (F, Cl, Br, and I) Group 18 the noble gases (He, Ne, Ar, Kr, Xe, Rn) 4
3. D-block elements (Groups 3-12) -- also called the transition elements -- e - configuration is somewhat less predictable for these elements (although most are normal) -- all have typical metallic properties -- generally have 2 valence e - (outer shell) 4. f-block elements (belong between groups 3 and 4; placed below the rest of the P.T.) usually have 2 valence e - Pass out lrg. P.T. QUIZ --Lanthanides found in Pd. 6 also called rare-earth elements although they aren t rare. Shiny, reactive metals. -- Actinides found in Pd. 7. Unstable and radioactive. The 1 st four are found naturally; the rest have only been made in the lab. III. Electron Configuration and Other Periodic Properties A. Atomic Radius distance from the nucleus to the edge of an atom. As atomic # increases across a period, more e - are added to the same shell, but more p are also added to the nucleus. This stronger + charge in the nucleus pulls the e - closer, decreasing the radius. As you go down a group, new e - are added to a new shell, farther from the nucleus, so radius increases. 5
B. Ionic Radius - Ionic sizes relative to other ions. Follows the same trends as atomic radius. Cation (cat-ion) a positively charged ion -- Formed by the loss of 1 or more e - -- Causes a decrease in radius compared to the neutral atom -- Most metals form cations (left side of table; low # of valence e -, easier to lose a few to get to octet than to gain several). Anion (an-ion) a negatively charged ion -- Formed by gaining 1 or more e - -- causes and increase in radius compared to the neutral atom -- Most nonmetals form anion (right side of table; high # of valence e -, easier to gain a few to get octet than lose several). C. D. Electronegativity the ability of an atom within a chemical compound dto attract electrons. -- Based on whether attracting e - gets them closer to an octet, and based on size (smaller atoms attract e - more because the e - are closer to the nucleus). Fluorine the most electronegative element. Assigned an arbitrary EN of 4.0 (all other elements have a lower EN) 6
In the compound HF, to which atom will the e - be closer? F How about the compound NaCl? Cl The compound H 2 O? O D. The Electron affinity - of an atom is the energy change when an electron is added to the neutral atom to form a negative ion. This property can only be measured in an atom in gaseous state. 7
E. Ionization Energy - the minimum energy required to remove (to infinity) an electron from the atom or molecule Summary of Periodic Trends Periodic Law Puzzle 8
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