Periodic Trends

Rohit C. Tilwani, RCh, PhD

 Periodic Trends

Learning Objectives:
Interpret group trends in atomic
radii, ionic radii, ionization energies,
electron affinity, and
electronegativity
Atomic Radius
Definition: Half of the distance between nuclei
in covalently bonded diatomic molecule

Radius decreases across a period

→ Increased effective nuclear charge due
to decreased shielding

Radius increases down a group

→ Each row on the periodic table adds a

“shell” or energy level to the atom
 Atomic Size

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Radius
Atomic Radius = half the distance between two
nuclei of a diatomic molecule.
 Trends in Atomic Size
Influenced by three factors:
1. Energy Level
– Higher energy level is further
away.
2. Charge on nucleus
– More charge pulls electrons in
closer.
3. Shielding effect e- <-> e- repulsion
 Group trends
H
• As we go down
Li
a group...
• each atom has Na
another energy
level, K

• so the atoms
get bigger. Rb
 Periodic Trends
• As you go across a period, the
radius gets smaller.
• Electrons are in same energy level.
• More nuclear charge.
• Outermost electrons are closer.

Na Mg Al Si P S Cl Ar
 Rb
K

Overall
Atomic Radius (nm)

Na

Li
Kr
Ar
Ne
H

10 Atomic Number
Periodic Trends
Table of
Atomic
Radii
 Atomic Radii
IA IIA IIIA IVA VA VIA VIIA

Li Be B C N O F
1.52 1.11 0.88 0.77 0.70 0.66 0.64

Na Mg Si
Al P S Cl

1.86 1.60 1.43 1.17 1.10 1.04 0.99

K Ca Ga Ge As Se Br

2.31 1.97 1.22 1.22 1.21 1.17 1.14

Rb Sr In Sn Sb Te I

2.44 2.15 1.62 1.40 1.41 1.37 1.33

Cs Ba Tl Pb Bi

2.62 2.17 1.71 1.75 1.46

= 1 Angstrom
Period Trend:
Atomic Radius
Ionization Energy
Definition: the energy required to remove an
electron from an atom

❑ Tends to increase across a period

❑ As radius decreases across a period,
the electron you are removing is closer to
the nucleus and harder to remove
❑ Tends to decrease down a group
❑ Outer electrons are farther from the
nucleus and easier to remove
 Ionization Energy
The second ionization energy is the
energy required to remove (1 mole
of) the second electron(s).
Always greater than first IE.
The third IE is the energy
required to remove a third
electron.
Greater than 1st or 2nd IE.
Symbol First Second Third
H 1312
He 2731 5247
Li 520 7297 11810
Be 900 1757 14840
B 800 2430 3569
C 1086 2352 4619
N 1402 2857 4577
O 1314 3391 5301
F 1681 3375 6045
Ne 2080 3963 6276
Symbol First Second Third
H 1312
He 2731 5247
Li 520 7297 11810
Be 900 1757 14840
B 800 2430 3569
C 1086 2352 4619
N 1402 2857 4577
O 1314 3391 5301
F 1681 3375 6045
Ne 2080 3963 6276
 What determines IE
• The greater the nuclear charge,
the greater IE.
• Greater distance from nucleus
decreases IE
• Filled and half-filled orbitals
have lower energy, so achieving
them is easier, lower IE.
• Shielding effect
 Shielding
• The electron in the
outermost energy
level experiences
more inter-electron
repulsion (shielding).
• Second electron has
same shielding, if it is
in the same period
 Group trends
• As you go down a group, first IE
decreases because...
• The electron is further away.
• More shielding.
 Periodic trends
• All the atoms in the same period
have the same energy level.
• Same shielding.
• But, increasing nuclear charge
• So IE generally increases from
left to right.
• Exceptions at full and 1/2 full
orbitals.
 He
• He has a greater
IE than H.
• same shielding
First Ionization

H
• greater nuclear
charge
energy

Atomic number
 He
Li has lower IE
than H
First Ionization

Outer electron
H further away
outweighs
energy

greater nuclear
Li
charge

Atomic number
 He
Be has higher IE
than Li
First Ionization

same shielding
H Be greater nuclear
charge
energy

Li

Atomic number
 He
B has lower IE
than Be
same shielding
First Ionization

H Be
greater nuclear
charge
energy

B p orbital is slightly
Li more diffuse and
its electron easier
to remove
Atomic number
 He
First Ionization

H C
Be
energy

B
Li

Atomic number
 He

N
First Ionization

H C
Be
energy

B
Li

Atomic number
 He
• Breaks the pattern,
N because the outer
electron is paired in a p
orbital and experiences
First Ionization

inter-electron repulsion.
H C O
Be
energy

B
Li

Atomic number
 He

N F
First Ionization

H C O
Be
energy

B
Li

Atomic number
 He Ne
• Ne has a lower
N F IE than He
• Both are full,
First Ionization

H Be
C O • Ne has more
shielding
energy

B • Greater
Li distance

Atomic number
 He Ne
Na has a lower
N F IE than Li
Both are s1
First Ionization

H C O Na has more
Be
shielding
energy

B Greater
Li
distance
Na

Atomic number
 First Ionization
energy

Atomic number
 Periodic Trend:
Ionization Energy
 Trends in Electron Affinity

The energy change associated with adding an

electron (mole of electrons) to a (mole of) gaseous
atom(s).
• Easiest to add to group 7A.
• Gets them to full energy level.
• Increase from left to right: atoms become
smaller, with greater nuclear charge.
• Decrease as we go down a group.
Electronegativity
Definition: A measure of the ability of an atom in
a chemical compound to attract electrons

o Electronegativity tends to increase

across a period
o As radius decreases, electrons get
closer to the bonding atom’s nucleus
o Electronegativity tends to decrease
down a group or remain the same
o As radius increases, electrons are
farther from the bonding atom’s
nucleus
Periodic Table of Electronegativities
 Periodic Trend:
Electronegativity
 Summary of
Periodic Trends
 Ionic Radii
❑ Positively charged ions formed when
an atom of a metal loses one or
Cations more electrons
❑ Smaller than the corresponding
atom

❑ Negatively charged ions formed

when nonmetallic atoms gain one
Anions or more electrons
❑ Larger than the corresponding
atom
 Trends in Ionic Size

• Cations form by losing electrons.

• Cations are smaller that the atom
they come from.
• Metals form cations.
• Cations of representative
elements have noble gas
configuration.
 Ionic size

• Anions form by gaining electrons.

• Anions are bigger that the atom
they come from.
• Nonmetals form anions.
• Anions of ‘A’ groups elements
have noble gas configuration.
 Configuration of Ions

• Ions have noble gas configurations

(not transition metals).
• Na is: 1s22s22p63s1
• Forms a 1+ ion: 1s22s22p6
• Same configuration as neon.
• Metals form ions with the
configuration of the noble gas
before them - they lose electrons.
 Configuration of Ions

• Non-metals form ions by gaining

electrons to achieve noble gas
configuration.
• They end up with the
configuration of the noble gas
after them.
 Group trends
• Adding energy
level Li1+
Na1+
• Ions get bigger as K1+
you go down.
Rb1
+

Cs1+
Graphic courtesy Wikimedia Commons user Popnose
 Periodic Trends
• Across the period, nuclear charge
increases so they get smaller.
• Energy level changes between
anions and cations.

N3-
B3+ O2- F1-
Li1+

Be2+ C4+
 Size of Isoelectronic ions

• Iso- means the same

• Iso electronic ions have the same
# of electrons
• Al3+ Mg2+ Na1+ Ne F1- O2- and N3-
• all have 10 electrons
• all have the configuration:
1s22s22p6
 Size of Isoelectronic ions
Positive ions that have more
protons would be smaller.

N3-
O2-
Ne F1-
Al3+ Na1+

Mg2+