section 5.3 – electron configuration and periodic properties
TRANSCRIPT
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HONORS CHEMISTRYSection 5.3 – Electron Configuration and Periodic Properties
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Atomic Radius
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
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Trends in Atomic Radii
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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)
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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?
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Ionization Energy
Atom + energy → Atom+ + e-
First electron removed – First Ionization Energy (IE1)
Second electron removed – Second Ionization Energy (IE2) 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
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Ionization Energies
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Ionization Energies
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Successive Ionization Energies
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Why?
Each successive electron feels a stronger nuclear attraction
This information lead to the understanding of the stability of the noble gas configuration+
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Practice
Choose the element with the higher IE1: Ca or Ba Ca or Br Ca or K Ca or Mg
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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!!!!
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Electron Affinity Values
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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
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Second Electron Affinities Very difficult to add an electron to an
anion (negative ion) Second Electron Affinities are all
positive
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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
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Ionic Radii
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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
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Bonded Atoms
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
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Electronegativity
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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
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Summary of Trends
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Summary of Trends
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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
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D and F Blocks
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