Download Chpater 5.3 PPT

Section 5.3 – Electron Configuration and Periodic Properties
HONORS CHEMISTRY
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
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
 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
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 IE1:
 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
 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
 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