Average values of the Duration
of Electronic Transitions
Kid Spoon
•Forward –18 s
•Backward –35 s
Nail
•Forward –4s
•Backward –28 s
Fever Strip
Forward -14 s
Backward –22 s
Ironing of color paper
Pink paper
Forward:
Instantaneous
Backward: 3.2 min
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electrons. It is also referred to as the screening effect (or) atomic shielding. The shielding effect also explains why valence-shell electrons are more.
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Average values of the Duration
of Electronic Transitions
Kid Spoon
- Forward – 18 s
- Backward – 35 s
Nail
- Forward – 4s
- Backward – 28 s
Fever Strip
Forward - 14 s
Backward – 22 s
Ironing of color paper
Pink paper
Forward:
Instantaneous
Backward: 3.2 min
Conclusions
• 1. All the four material investigated are
Thermoluminescent materials. They
absorb heat and emit as light.
• 2. The backward color change is always
slower than the forward color change
• The backward color change involves
several phenomena whereas the forward
is single and straightforward.
Periodic Trend in Ionic Radii: The Group Ionic radius increases as we move down the group. The reasons for this trend are: (1) Number of shells increases; (2) Shielding also increases Periodic Trend in Ionic Radii: The Period Ionic radius decreases as we move from Left to Right in a Period. The reasons for this trend is: electrons are added to the same shell; therefore, nucleus attracts them continuously.
Shielding of the Valence Electron by the Inner Electrons
The shielding effect is the reduction
in the effective nuclear charge on the
outer electrons, due to the inner
electrons. It is also referred to as
the screening effect (or) atomic
shielding.
The shielding effect also explains
why valence-shell electrons are more
easily removed from the atom.
The plot of atomic number versus atomic radius gives a measure of the shielding effect.
The Shielding Effect can be defined as: Diminishing of the force or the control of the
nucleus on outer electrons by the inner electrons, which act as a curtain between the
nucleus and the outer electrons.
The graph above represents the change in ionic radius with atomic number. Ionic Radius is the distance between the center of the nucleus and the valence shell. In both graphs, we see that there is decrease in ionic radius with atomic number. The left graph is for Lanthanides, which are “f” block elements. The right graph is for fourth period elements, which include Calcium (which is a “s” block element) and Sc through Zn (which are “d” block elements called Transition Elements). For the lanthanides, there is a dramatic fall in the ionic radius whereas for the fourth period elements, the decrease is not steep. The steep decrease for lanthanides is due to lanthanide contraction, which is reduction in the size of lanthanide ions. In lanthanides, electrons are continuously added to the “ 4 f” orbitals, which are in the 4 th^ shell much inner to the valence shell, which is the 6 th^ shell. For the fourth period elements, there is only a small fall in the ionic radius, this is because after Calcium, electrons are added to the “ 3 d” orbitals, which are in the 3 rd^ shell just inner to the valence shell, which is the 6 th^ shell.
https://www.thestudentroom.co.uk/showthread.php?t= 3817571
A is the correct answer as anion is bigger than the cation and more the charge, the
smaller is the cation and the bigger is the anion.
Ionic radius decrease across periods because effective nuclear charge increases. That is, the net positive charge experienced by an electron in the ion increases as a result of the number of protons in the nucleus increasing.
Ionization Energy in general is the energy that must be supplied to an atom to form
cations. Ionization Energy can be First Ionization Energy or Second Ionization Energy or
Third Ionization Energy corresponding successively to the formation of Unipositive Cation
from atom; Dipositive cation from Unipositive Cation; and Tripositive cation from
Dipositive cation.
First Ionization Energy {Atom to Unipositive Cation} The amount of energy required to completely remove an electron from a gaseous atom to form a unipositive cation. X(g) + IE 1 →X
- e-
The second ionization energy {Unipositive Cation to Dipositive Cation}
- X+^ (g) + IE 2 X2+^ (g) + e-
The third ionization energy {Dipositive Cation to Tripositive Cation}
- X2+^ (g) + IE 3 X3+^ (g) + e-
More energy required to remove 2
nd
electron, and still more energy required to
remove 3rd^ electron
Generally, IE3 > IE2 > IE 1
IONIZATION ENERGY
Why does the value of IE 1
decrease within a group and
increase within a period?
TRENDS
- The ionization energy of the elements within a period generally increases
from left to right. This is due to valence shell stability.
- The ionization energy of the elements within a group generally
decreases from top to bottom. This is due to electron shielding.
Electron Affinity The apex points are occupied by halogens; they have the highest electron affinity. This is because halogens are the most electronegative elements as they are next to the Noble Gases and are in need of just one electron to become the Noble Gases. The bottom points are occupied by (i) Noble Gases (ii) Alkaline Earth Metals (iii) Mg, Mn, Zn, Cd, Hf, and Hg, which are stable transition metals (iv) Nitrogen Noble gases have fully filled valence shell; so they do not need electrons and therefore they have very low electron affinity. Alkaline earth metals and the stable transition elements; they are electropositive and do not need electrons; therefore, they have very low electron affinity. Nitrogen is a non-metal and it is expected to have higher Electron Affinity as non-metals are electronegative. Nitrogen although a non-metal has low Electron Affinity; this can be explained as follows: