Section 5.3 – Electron Configuration and Periodic Properties
HONORS CHEMISTRY

Atomic Radius
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Determined by the distance from the nucleus to the edge of the outer orbital.
Edge of outer orbital not well defined
Use identical bonded atoms – ½ of the distance between the nuclei

Trends in Atomic Radii

Trends in Atomic Radii


Decrease across the period
 Due to increasing positive charge of the nucleus
Increase as you go down the group
 Exception Ga to Al – Ga smaller due to increased nuclear charge (first addition of d electrons)

Problems
 Which of these elements; Li, Rb, K or Na has the smallest radius? Largest?
 Which of these elements; O, Se, S and Po has the smallest atomic radius? Largest?

Ionization Energy
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Atom + energy → Atom + + e -
First electron removed – First Ionization Energy
(IE
1
)
Second electron removed – Second Ionization
Energy (IE
2
) etc.
Group 1 – lowest ionization energy
Group 18 - highest ionization energy
Ionization energies increase across the period due to increased nuclear charge.
Ionization energies decrease down a group due to further distance from nucleus and electron shielding

Ionization Energies

Ionization Energies

Successive Ionization Energies

Why?


Each successive electron feels a stronger nuclear attraction
This information lead to the understanding of the stability of the noble gas configuration+

Practice
 Choose the element with the higher IE
1
:
 Ca or Ba
 Ca or Br
 Ca or K
 Ca or Mg

Electron Affinity
 Atom + e → Atom + (- energy)
 IMPORTANT!!!!!
 Negative energy means energy lost by system
 Positive energy means energy gained by system
 Sign indicates direction not numerical value!!!!

Electron Affinity Values

Electron Affinity Trends


Generally become larger (look at as absolute value) as you move across the period.
Exception – Group 15 due to half filled p orbitals
 Generally become smaller as you move down a group due to:
 Greater Nuclear Attraction
 Greater Atomic Radius

Second Electron Affinities


Very difficult to add an electron to an anion
(negative ion)
Second Electron Affinities are all positive

Ionic Radii


Cation (positive ion)
 Smaller atomic radius than atom
 Due to:
 electrons being removed
 increased effective nuclear charge
Anion
 Larger atomic radius than atom
 Due to
 electrons being added
 decreased effective nuclear charge
 Greater repulsion of electrons

Ionic Radii

Valence Electrons



Available to be lost, gained or shared when compounds are formed.
In outer main energy levels
For Main Group Elements – s and p orbitals

Bonded Atoms
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Very rarely are electrons shared equally
Usually attracted more to one atom

This will effect the chemical properties of the compound!!!
Measure of attraction – called electronegativity
Based on a 4.0 scale – F = 4.0.
Developed by Linus Pauling

Electronegativity

Electronegativity Trends


Increase across a period.
Tend to decrease or stay the same down a group.


If a noble gas does not form compounds – it does not have an electronegativity
If a noble gas does form compounds – it will have a high electronegativity

Summary of Trends

Summary of Trends

D-Block


These elements tend to vary less and with less regularity than Main Group Elements.
Still electrons in d orbitals are often responsible for characteristics of elements in the d-block


Atomic radius tends to decrease across the block
Ionization energies generally increase across both the d and f-blocks

D and F Blocks
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Tend to lose electrons from outer shell!!!!
That means the valence electrons come from the ns shell not the (n-1)d shell
Generally these elements from 2+ ions.
Electronegativities
 D-block - between 1.1 and 2.54 (Only groups 1 and
2 are lower)
 F-block – between 1.1 and 1.5