The document discusses several periodic trends: 1. Atomic radius generally increases moving down a group and decreases moving across a period, due to shielding and nuclear charge effects. 2. Ionization energy increases moving across a period as the effective nuclear charge increases, and decreases moving down a group as the distance from the nucleus increases. 3. Electronegativity increases moving across a period as it is easier to gain than lose electrons, and decreases moving down a group with increased atomic radius.

1. Course: B.Sc.
Microbiology/Biotechnology/Biochemistry
Sem II
Sub: Inorganic Chemistry
Unit 2
Periodic properties

2. Atomic RadiusAtomic Radius
Radius
Atomic radius is the
distance from the center of
an atom’s nucleus to its
outermost electron.

3. Atomic Radii
• The atomic radius : size of the atom.
• determine : finding half the length between
the nucleus’ of two atoms of the same
elements.

4. AtomicRadiusincreases
Atomic Radius increases
H
Li
Na
K
Fr
F
At
Atomic RadiusAtomic Radius

5. Atomic Radii
0
0.05
0.1
0.15
0.2
0.25
1 2 3 4 5 6 7 8
Element (atomic number)
Radius(nm)
Series1

6. Atomic RadiusAtomic Radius
Trend #1Trend #1
Atoms get
larger going
down a group
Hydrogen
(1s1
)
Lithium (1s2
2s1
)
Sodium (1s2
2s2
2p6
3s1
)
Shielding EffectShielding Effect

7. Atomic RadiusAtomic Radius
Trend #2Trend #2
Shrinking EffectShrinking Effect
5p+ 6p+
7p+
1s2
2s2
2p1
1s2
2s2
2p2
1s2
2s2
2p3

8. Trends within periods
• The size of atoms (and therefore the atomic radii) increase as we go down groups
on the periodic table (since more energy levels are added), BUT as we go to the
right, the size decreases because…
• Electrons are being added to the same shell,
which does not change the size of the atom,
but protons are also being added to the
nucleus, which pulls the electrons inwards
towards the center, decreasing the radius.

9. Look at the trend within the periods.

10. Ionic Radius
• The ionic radius is the radius of an atom’s ion.
• When an atom gains or loses electrons, it
becomes an ion.
• Because an electron is negatively charged,
atoms that lose or gain acquire a net charge
– If an atom loses electrons it becomes _________
– If an atom gains electrons it becomes _________

11. Ionic Radius Continued
• atom - loses electrons : become smaller.
• nucleus : able to pull them closer.

12. • atoms - gain electrons : become larger.

13. • Silicon4+:
(the atoms lose more and more electrons,
making the pull from the nucleus stronger each
time).
• Silicon4-
: the atomic radius increases by a lot,
because it has 4 extra electrons that repel, forcing
the ion out (and therefore getting bigger). Then the
ionic radii continues on it’s decreasing pattern,
because the ions after silicon each have one less
extra electron then the previous element, and
therefore push out (repel) less.

14. Look at the trends for ionic size as you
move across a period.

15. Ionization EnergyIonization Energy
Energy needed to remove one of atom’s electrons from its
outermost shell
A + E A+
+ e-
Reflection of how strongly an atom holds onto
its outermost electron.
Atoms with high ionization energies hold ontoAtoms with high ionization energies hold onto
their electrons very tightly.their electrons very tightly.
Atoms with low ionization energies are moreAtoms with low ionization energies are more
likely to lose one or more of their outermostlikely to lose one or more of their outermost
electron.electron.

16. Ionizationenergyincreases
Ionization energy increases
H
Li
Na
K
Fr
F
At
Ionization EnergyIonization Energy

17. • first ionization energy : from neutral atoms to cations with a
1+ charge.
• second ionization energy : form 2+ cations from 1+ cations:
• M+
(g) M2+
(g) + e-
• third ionization energy : form 3+ cations:
• M2+
(g) M3+
(g) + e-
- IE : positive numbers because energy must be supplied (an
endothermic energy change) to separate electrons from
atoms. .

18. Periodic table:
IE : decreases from top to bottom in groups, and increases from
left to right across a period.
He : largest first ionization energy
francium : one of the lowest.
• From top to bottom :.
• orbitals - higher values of the principal quantum number (n),
which : further away from the nucleus.
• Since the outermost electrons are further away, they are less
strongly attracted by the nucleus, and are easier to remove,
corresponding to a lower value for the first ionization energy.

19. • From left to right : more protons are being added to the
nucleus, but the number of electrons in the inner, lower-
energy shells remains the same.
• The valence electrons feel a higher effective nuclear charge —
the sum of the charges on the protons in the nucleus and the
charges on the inner, core electrons.
• The valence electrons are therefore held more tightly, the
atom decreases in size , and it becomes increasingly difficult
to remove them, corresponding to a higher value for the first
ionization energy.
•

20. Metallicity
Ability of an atom to lose an electron
TREND:
Increases from top to bottom
Decreases from left to right

21. Electron AffinityElectron Affinity
A + e-
A-
+ E
Measure of an atom’s attraction, or affinity, for
an extra electron.
Energy released when an atom gains an
electron to form a negative ion/anion.
increases within a period from left to right.
- As one goes down a group, electron affinity
decreases.

23. • Periodic Trends :
• increases : left to right across the periodic table, from the
alkali metals to the halogens.
• small changes in the electron affinity are observed as you
move down a group.
• electron affinity : measure of how stable the products are
with respect to the reactants.
• products - much more stable, a large amount of energy will
be released during the process and EA will be a large negative
number.
• reactants - much more stable than the products, then it
becomes very difficult to add an electron and the EA will be
positive.

24. • Halogens (group 7A, F to At) : Most negative EA values,
addition of an e-
leads to noble gas configuration, very
favorable.
• Group 5A (N to Bi): filled shell discourages addition of an
electron, EA values less negative than neighbors (groups 4A &
6A).
• Alkaline Earths (group 2A, Be to Ba) : Filled s-subshell
discourages addition of an electron, EA values nearly zero.
• Noble Gases (Group 8A, He to Rn) : Completely filled shell
strongly discourages addition of an electron, EA values are
positive.

25. ElectronegativityElectronegativity
Ability of an atom to attract electrons in a
chemical bond.
Electronegativityincreases
Electronegativity increases
H
Li
Na
K
Fr
F
At

26. Periodic table:
• Move to the right : electronegativity increases.
• valence shell of an atom is less than half full -less energy to
lose an electron than gain one and thus, it is easier to lose an
electron.
• Conversely, when the valence shell is more than half full, it is
easier to pull an electron into the valence shell than to donate
one.

27. • Move down a group: decreases.
• because the atomic number increases down a group and thus
there is an increased distance between the valence electrons
and nucleus, or a greater atomic radius.
• Important exceptions :noble gases, lanthanides, and
actinides.
• noble gases : complete valence shell .
• lanthanides and actinides : more complicated chemistry that
does not generally follow any trends.

28. • Therefore, noble gases, lanthanides, and actinides :
electronegativity values.
• for transition metals : while they have values, there is little
variance among them as you move across the period and up
and down a group. This is because of their metallic properties
that affect their ability to attract electrons as easily as the
other elements.

29. Electronegativityincreases
Electronaffinityincreases
Electronegativity increases
H
Li
Na
K
Fr
F
At
Ionizationenergyincreases
Ionization energy increases
Electron Affinity increases
Summary of PeriodicSummary of Periodic
TrendsTrends
AtomicRadiusincreases/ionic
size
Metallicity inreases
Atomic Radius increases/ Ionic size
Metallicityincreases

30. • Summary of Periodic Table Trends
• Moving Left → Right
• Atomic Radius Decreases
• Ionization Energy Increases
• Electronegativity Increases
• Moving Top → Bottom
• Atomic Radius Increases
• Ionization Energy Decreases
• Electronegativity Decreases

31. Look at the trends in ionic size as you
move down a group

32. References
• Essentials of Physical chemistry by Bahl Arun, S Chand, 2012
• General chemistry by Ebbing Darrell D, 5th
, A I T B S Publishers,
2002