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 PropertiesAtomic RadiiOne-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 energiesAlways 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

ElectronegativityMeasure of the ability of an atom to attract electrons in a chemical bond

F – highest electronegatvity! 4.0

Period TrendsE.N. tends to increase across each period

Are some exceptions – don’t worry about those!

Group trendsE. N. tend to either decrease down a group or remain about the same.

Periodic properties of the d and f block elementsNot holding you responsible for these

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