History of the Periodic Table

1. History of the Periodic Table Russian scientist Mendeleev determined that when atoms are arranged in order of increasing atomic mass chemical properties appeared at regular intervals. Mendeleev accurately predicted the existence and properties of three elements that had not been discovered yet. Sc, Ga, Ge Mendeleev was credited with discovering the periodic Law: The physical and chemical properties of the elements are periodic functions of their atomic number. In 1911 English Scientist Henry Mosely arranged the elements by increased atomic number, which is the basis for today’s periodic table.

2. Trends in Some Periodic Properties The physical and chemical behavior of the elements is based on the electron configurations of their atoms. e- configurations can be used to explain many of the repeating or “periodic” properties of the elements

3. Atomic Radius The atomic radius decreases across a period This is because of an increase in positive charge in the nucleus because of an increase in the number of protons. The atomic radius increases down a group This is because the electrons occupy sublevels in higher energy levels which increases the distance between the nucleus and the electron.

4. Electron Shells and Sizes of Atoms • Size in main group elements follow 2 general rules • 1. Atomic Radius increases going down in a group • 2. Atomic Radius decreases going left to right in a period Atomic radii of the main-group and transition elements.

5. PROBLEM: Using only the periodic table (not Figure 8.15), rank each set of main group elements in order of decreasing atomic size: SAMPLE PROBLEM 8.3 Ranking Elements by Atomic Size (a) Ca, Mg, Sr (b) K, Ga, Ca (c) Br, Rb, Kr (d) Sr, Ca, Rb PLAN: Elements in the same group increase in size and you go down; elements decrease in size as you go across a period. SOLUTION: (a) Sr > Ca > Mg These elements are in Group 2A(2). (b) K > Ca > Ga These elements are in Period 4. (c) Rb > Br > Kr Rb has a higher energy level and is far to the left. Br is to the left of Kr. (d) Rb > Sr > Ca Ca is one energy level smaller than Rb and Sr. Rb is to the left of Sr.

6. Ionization energy defn:Ionization Energy: The energy required to remove one electron from a neutral atom. In general the ionization energy increases across a period. This is because of the increase attraction of the nucleus across a period Ionization energy generally decreases down a group. This is because of the increase in distance between the nucleus and the electron.

7. PROBLEM: Using the periodic table only, rank the elements in each of the following sets in order of decreasing IE1: SAMPLE PROBLEM 8.4 Ranking Elements by First Ionization Energy (a) Kr, He, Ar (b) Sb, Te, Sn (c) K, Ca, Rb (d) I, Xe, Cs PLAN: IE decreases as you proceed down in a group; IE increases as you go across a period. SOLUTION: (a) He > Ar > Kr Group 8A(18) - IE decreases down a group. (b) Te > Sb > Sn Period 5 elements - IE increases across a period. (c) Ca > K > Rb Ca is to the right of K; Rb is below K. (d) Xe > I > Cs I is to the left of Xe; Cs is furtther to the left and down one period.

8. Electronegativity defn: Electronegativity: A measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound. In general Electronegativity increases across a period In general Electronegativity decreases down a group Fluorine is the most electronegative element

9. defn: Cation: a positively charged ion defn: Anion: A negatively charged ion.

10. Figure 8.21 Trends in three atomic properties