honors chemistry chapter 6 electron configuration and the periodic table relationship between period...
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Periodic Table continuedHonors Chemistry Chapter 6
Electron configuration and the Periodic TableRelationship between period length and sublevels being filled
“Blocks” on the table – be able to identify
s,p,d,f
Group 1All have ns1 outer shell notation
Group 2All have ns2 outer shell notationThe value for n tells you what period
it is in, the superscript lets you know the group
d block elementsGroups 3 – 12
(n-1)dns
Add together the outermost d and s electrons and it will equal the group number
p blockGroups 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 blockLanthanide seriesActinide seriesf sublevel being filledLanthanide series – shiny metals similar in
reactivity to Group 2 – alkaline earth metalsActinide 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
TRENDS IN ATOMIC RADII 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)
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
ionization energy
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- + energyQuantity 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
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