Chapter 5 The Periodic Law

Sect. 5-1: History of the Periodic Table Stanislao Cannizzaro (1860) proposed method for measuring atomic mass at First International Congress of Chemists Dmitri Mendeleev (1869) arranged elements by atomic mass & similar chemical properties; left blanks for undiscovered elements

Henry Moseley (1911) arranged periodic table by atomic number Periodic law – properties of elements are periodic functions of their atomic #’s Noble gas group, lanthanide and actinide series added later

Sect. 5-2: Electron Configuration and the Periodic Table S-block elements are highly reactive metals because they easily give up their 1 or 2 valence electrons Group 1 – Alkali metals Silvery, can be cut with knife Group 2 – Alkaline Earth metals Harder, denser, stronger, and slightly less reactive than group 1

Special cases: Hydrogen grouped with 1 because of electron configuration, but doesn’t share their properties Helium is grouped with 18 because it has similar properties since its outside energy level is full, even though it has the same electron configuration as group 2

D-block elements Total # electrons in d plus electrons in highest s orbital = group # Referred to as transition elements Good conductors of heat/electricity High luster Not as reactive as s-block elements

P-block elements Combined with s-block they are called main-group elements Contains metals, non-metals, and metalloids, thus wide range of properties Group 17 – halogens Most reactive nonmetals Group 18 – noble gases nonreactive

F-block elements Lanthanides Shiny, similar in reactivity to group 2 Actinides All radioactive First 4 have been found naturally, all others are man- made

Sect. 5-3: Electron Configuration and Periodic Properties Atomic Radius – one half the distance between the nuclei of chemically bonded identical atoms Decreases from left to right across a period due to higher positive charge on right pulling electrons closer Increases going down a group because of adding energy levels

Ion – charged particle Ionization energy (IE) – energy required to remove an electron from a neutral atom in the gas phase Increases as you move to the right because those elements will less readily give up an electron Decrease as you move down a group due to electrons being further away from nucleus and shielded by inside electrons

2 nd and 3 rd ionization energies refer to removing additional electrons from positively charged ions 2 nd and 3 rd Ionization energies have a drastic “jump” if the ion has the electron configuration of a noble gas

Electron Affinity – energy change when a neutral atom gains an electron Reported as a negative # because of loss of energy Generally decreases as you move down a group Generally decreases as you move left on a period Exceptions for half-filled or filled sublevels Adding additional electrons will always have a positive value (requires energy)

Ionic Radii Cation – positively charged ion (lost electron) Will decrease radius because of loss of outer energy level Anion – negatively charged ion (gained electron) Will increase radius because protons “pulling in” are the same and with extra electrons they repel each other and spread out Cation & anion radius increases from the right to the left across a period Cation & anion radii increase down a group

Valence Electrons – electrons in outermost energy level (can be gained lost or shared) For s-block, # valence electrons is equal to group number For p-block, # valence electrons is equal to group number minus 10

Electronegativity – measure of ability of an atom in a compound to attract electrons Generally decrease as you move to the left of a period Generally decrease or stay the same moving down a group Nitrogen, Oxygen, and Halogens are most electronegative

Trends for d- and f-blocks atomic radius trend is same a main group, but with smaller changes Ionization energy trend is same for period, but increases going down a group Ion formation – electrons are removed from the s orbital 1 st, then the d Electronegativity trends are same