Periodic Table continued Honors Chemistry Chapter 6
Electron configuration and the Periodic Table • Relationship between period length and sublevels being filled • “Blocks” on the table – be able to identify • s,p,d,f
Group 1 • All have ns1 outer shell notation • Group 2 • All have ns2 outer shell notation • The value for n tells you what period it is in, the superscript lets you know the group
d block elements • Groups 3 – 12 • (n-1)dns • Add together the outermost d and s electrons and it will equal the group number
p block • Groups 13 – 18 • (with groups 1 and 2 are called the “main group” or “representative” element) • general electron configuration for p block is ns2np • Metals, metalloids, and nonmetals contained in this block.
f block • Lanthanide series • Actinide series • f sublevel being filled • Lanthanide series – shiny metals similar in reactivity to Group 2 – alkaline earth metals • Actinide series – all radioactive. Thorium through neptunium are found naturally on Earth. Others are laboratory made.
Periodic Properties • Atomic Radii • One-half the distance between the nuclei of identical atom s that are bonded together
Gradual decrease as atomic number increases across a period Caused by the increasing positive charge of the nucleus In general, atomic radii of the main group elements increases down a group (as a.n. increases) Trends in atomic radii
Ionic radii • Radius resulting when an atom forms an ion • Cation – positive ion. Results when a neutral atom loses electrons. Radius decreases • Anion – negative ion. Results when a neutral atom gains electrons. Radius increases
Ionization energy • The energy required to remove one electron from a neutral atom of an element (first ionization energy) • A + energy A+ + e- • Forms an “ion” – atom or group of bonded atoms that has a positive or negative charge • Process called “ionization” • Pg. 143 Table of ionization energy
Period trends • In general, first ionization energies increase as atomic number increases across a period for main-group elements • Metals – lose their electrons easily (reason for high reactivity) • Noble gases – highest i.e. values. Do not lose electrons easily – (accounts for low reactivity) • Increased nuclear charge accounts for increase in i.e.
Group trends • Among the main-group elements, i.e. generally decreases down the groups • Removed more easily because they are in higher energy levels, farther from the nucleus – able to overcome nuclear charge
2nd and 3rd ionization energies • Always higher than the first
Electron affinity • The energy change that occurs when an electron is acquired by a neutral atom • Most atoms release energy when this happens • A + e- A- + energy • Quantity of energy represented by a negative number
Some atoms must be “forced” • A + e- + energy A- • this quantity represented by a positive number • Ion made this way is very unstable – will lose the added electron spontaneously
Period trends • Halogens gain electrons most readily – reason for high reactivity • In general, as electrons are added to the same p sublevel with the same period, electron affinities become more negative • There are exceptions to this
Group trends • Not as regular as trends for i.e. • As a general rule, electrons add with greater difficulty down a group
Valence electrons Are the electrons available to be lost, gained, or shared in the formation of chemical bonds Often located in incompletely filled main-energy levels
Electronegativity • Measure of the ability of an atom to attract electrons in a chemical bond • F – highest electronegatvity! 4.0
Period Trends • E.N. tends to increase across each period • Are some exceptions – don’t worry about those!
Group trends • E. N. tend to either decrease down a group or remain about the same.
Periodic properties of the d and f block elements • Not holding you responsible for these
Yay!!! • Fire water • periodic trends tutorial 1 • periodic trends tutorial 2