Chemistry – [Periodicity & Group Trends] 1
Periodicity:
1. Terminology:
- Nuclear Charge
o Charge of the all the protons in the nucleus.
- Effective Nuclear Charge
o Effective nuclear charge = net positive charge felt by the valence electrons, after
deducting the shielding effect by the inner electrons.
2. Factors Affecting Ionization Energy & Electron Affinity & Electrostatic Attraction:
- Distance of the valence electron from the nucleus:
o The further the valence electron is from the nucleus, the weaker the attraction pull
towards the nucleus.
o Ionization energy will decrease.
- Nuclear charge:
o As the number of protons in the nucleus increases, the nuclear charge will also
increase (more positive).
o Its attraction on the valence electrons gets stronger.
o Ionization energy will increase.
- Shielding effect by inner energy level electrons:
o Electrons in the inner energy level will not allow the nuclear charge to pass through;
as a result, the valence electron experiences a lower nuclear charge than the inner
electrons.
o Electrons in shells of lower energy levels are more effective than those in higher
energy levels in shielding.
3. PERIODIC TRENDS:
Trends Across Period
Decrease across period
Atomic Radius
- More electrons are added are added to the same shell
- More protons are added to the nucleus
- The shielding effect is almost constant for all the elements across a
period
- Electrons being pulled closer towards nucleus by increasing +ve
charge and increasing effective nuclear charge
Cations (+ve)
Decrease across period
Ionic Radius
Smaller than their corresponding atomic radii
- Electrons are removed in the formation of cations.
- Stronger pull on remaining electrons towards the
positively charged nucleus.
Anions (-ve)
Decrease across period
Greater than their corresponding atomic radii
- Electrons are added to the outer energy level,
forming an anion.
- Attraction of the electrons is weaker.
- Repulsion between the negatively charged clouds.
Ionization Energy
General increase across the period due to:
- Increasing effective nuclear charge
- Decreasing atomic radius
- Constant shielding effect
Electron Affinity
Across a Period, electron affinity increases due to:
- Atomic radius decreases
 The electron is added to energy level closer to the nucleus.
 Stronger attraction between the nucleus and the electron.
Increases across the group are there is an increase in the number of electrons.
Effective Nuclear
Charge
Trends Down Group
Atomic Radius
Increase down group
- More electron shells are added due to increasing n number
- Distance between the electrons and the nucleus increase as well
- Nuclear charge increases as more protons are added.
- Shielding effect by inner electrons increases.
- Effective nuclear charge remains constant.
- Strength of attraction of the electrons to the nucleus is not as strong.
Ionic Radius
Increase down group due to the addition of electron shells.
Ionization Energy Decreases down group due to:
Chemistry – [Periodicity & Group Trends] 3
Electron Affinity
Effective Nuclear
Charge
- Increase in the atomic radius
- Increase in shielding effect
- Effective nuclear charge remains constant
Down a Group, electron affinity decreases due to:
- Increase in Atomic radius
 The electron is placed in a higher energy level, further from the
nucleus.
 Attraction between the electron and nucleus decreases.
Remains around the same down the group.
(DON'T USE THIS TO EXPLAIN COS IT MIGHT BE DECREASING OR WHATEVER SO
YAH ON THE SAFE SIDE KEEP OFF THIS)
Summary of Trends:
Commonly Asked PPA Questions for Periodicity:
What to do: Use general trends to help you answer questions on specific elements!
Underline = variable, question can change it, but answering to the question is still largely similar.
Atomic Radius:
Explain why atomic radius increases down a group. (Explain general trends)
- Down a group, there is an increase in the number of electron shells. Therefore, the valence
electrons are further away from the nucleus as you go down a group.
Explain why the atomic radius of potassium is bigger than that of sodium. (Compare size of specific
elements in the same group/period)
- K has one more electron shell than sodium, and hence the valence electrons of K are in the
4s orbital which is further away from the nucleus as compared to Na’s valence electron in
the 3s orbital.
Ionization Energy:
Explain why IE increases across the period. (Explain general trends)
- Across the period, the effective nuclear charge increases, as the number of electron shells
remains the same while the nuclear charge increases. Hence, the valence electrons are more
strongly attracted to the nucleus, making them require more energy to remove.
Why is the 1st IE of oxygen greater than that of the 1st IE of lithium? (Compare 1st IE of elements in
same group)
- O has a greater effective nuclear charge as compared to Li, hence the valence electrons of O
are more strongly attracted to the nucleus, making them require more energy to remove
than Li.
Explain why IE increases with every successive electron removed. (Explain general trend of
successive IE)
- Each successful removal of an electron from an atom results in an increasingly positive ion,
hence the remaining valence electrons are more strongly attracted to the nucleus, making
them require increasingly more energy to remove.
Why is the 2nd IE of magnesium greater than that of its 1st IE? (Comparison between 2 IE values that
are not influenced by inner shell)
- The remaining valence electrons of Mg after its 1st IE are more strongly attracted to the
nucleus, making them require more energy to remove.
Why is there a big difference between the 1st IE of sodium and the 2nd IE? (Removal of electron from
inner shell)
- In Na the 1st electron is removed from the 3s valence orbital. However, the 2nd electron is
removed from the 2p orbital, which is part of the inner shell, hence requiring much more
energy to remove.
Ionic Radius/Cations/Anions:
Why is the ionic radius of Ca2+ smaller than that of Cl-? (Why ionic radius of cations < anions)
- Both ions have the same number of electrons, but Ca2+ has 20 protons while Cl- has 17
protons. Therefore, the electrons in Ca2+ experience a stronger attraction to the nucleus
and the ion is smaller.
Why is the ionic radius of Mg2+ smaller than that of Ca2+? (Compare the ionic radius of cations/anions
of same charge)
- Ca2+ has 3 electron shells, 1 more than the 2 electron shells Mg2+ has. Hence, the Mg2+
electrons are closer to the nucleus and hence are more strongly attracted to the nucleus
than Ca2+, making the ion smaller.
Halogens/Salt Formers (Group VII, Group 17)
1. General Properties & Trends
Physical Properties
Character
Non-metal, p-block elements
Density
Usually have a low density
Melting point
Down the group from F to At, the melting point of the elements increases,
with the physical state changing from gas  solid
Chemistry – [Periodicity & Group Trends] 5
Chemical Properties
Electron
Configuration
Electron Affinity
Reactivity
F>Cl>Br>I>At
Have 7 electrons in their valence shell
Electronic Configuration: ends with ns2np5
High EA as they are only one electron short of gaining stability.
Reactivity of the halogens DECREASES down the Group:
1. Atomic radius increases with higher principal quantum number due to the
addition of more electron shells.
2. The ‘incoming’ electron is increasingly shielded from the positively charged
nucleus by inner core electrons. (Shielding effect increases)
3.Weaker electrostatic force of attraction to the nucleus to form a halide ion.
2. Elements in Group VII
Element
F
(Fluorine)
Cl
(Chlorine)
Br
(Bromine)
I
(Iodine)
At
(Astatine)
Physical
State
Pale yellow
gas
Colourless
Pale greenyellow gas
Colourless
Reddishbrown
liquid
Yellow
liquid
(Dilute in
aq form)
Black Solid
Brown
Liquid (In
solution)
Purple Gas
?
Compound Colour formed
with K (Potassium)
(NIL)
Hardly see this existing)
Melting Point (K)
Density (g/cm3)
54
0.001696
Colourless
172
0.003
Colourless
266
3.103
Colourless
387
4.94
NIL
575
~7.00
* No need know exact values but 
3. Displacement Reactions
Halogen
added
Observation
Chlorine
Appearance of
aqueous solution
Potassium
chloride
Colourless
Potassium
bromide
Orange
Potassium iodide
Brown
Bromine
Iodine
Appearance of
cyclohexane
solution
Colourless
Reddish-brown
(Caused by
bromide)
Violet
(Caused by iodide)
Displacement?
No reaction
Cl2 + 2KBr 
2KCl + Br2
Cl2 + 2KI 
2KCl + I2
Appearance of
aqueous solution
Orange
Orange
Brown
Appearance of
cyclohexane
solution
Reddishbrown
Reddish-brown
Violet
Displacement?
No reaction
No reaction
Br2 + 2KI  2KBr + I2
Appearance of
aqueous solution
Brown
Brown
Brown
Appearance of
cyclohexane
solution
Violet
Violet
Violet
Displacement?
No reaction
No reaction
No reaction
What are displacement reactions?
- Displacement reactions are chemical reactions used to compare the reactivity of
elements.
- They are chemical reactions in which a less reactive element is replaced/displaced in a
compound by a more reactive one.
- Displacement reactions can occur either in Group VII (Halogens) or metallic elements.
What is cyclohexane? (*I don’t think we need to know but just in case)
It is a non-polar solvent with the molecular formula of C6H12
- Why different colourations: To do with polarity
Alkaline Metals (Group 1, Group I)
1. General Properties + Trends
Physical Properties
Very soft metal, s-block elements, stored in oil
Character
Have a lower density than other metals. Hence, Li, Na and K tend to float on
Density
water.
Down the group from Li to Fr, the melting point of the elements decreases.
Melting point
Generally, their melting points are quite low compared to other metals.
Chemical Properties
Has 1 electron in their valence shell
Electron
Chemistry – [Periodicity & Group Trends] 7
Configuration
Ionization Energy
Reactivity
Fr>Cs>Rb>K>Na>Li
Electronic Configuration: ends with ns1
Low ionization energy, can easily lose 1 electron
1.They are the most reactive metals, with low 1st ionization energy.
2.They are never found as free elements in nature – they are always bonded
with other elements to form compounds.
3. They react violently with water to form alkaline solutions.
4. Reactivity increases down a group as ionization energy decreases.
 Atomic radius Increases
 Shielding effect Increases (+electron shells)
 Force of attraction between protons and valence electron
Decreases
 1st ionization energy Decreases
2. Elements in Group I
Element
Li (Lithium)
Melting Point (K)
453.7
Density (g/cm3)
0.535
Na
371.0
(Sodium)
K
336.8
(Potassium)
0.971
Rb
(Rubidium)
312.2
1.5
Cs
(Caesium)
301.6 (Able to
melt in your
hands since body
temp is 310K)
300
1.87
Fr
(Francium)
0.862
2.4
3. Reactions with Water
Chemical Formula:
2X (s) + 2H2O (l)  2XOH (aq) + H2 (g), where X is any Group I metal.
Reaction Description when placed in water:
- Metal darts around surface (sometimes even becoming a ball) and fizzes. Effervescence is
observed.
- Hydrogen gas is liberated (produced). (Use lighted splint to test, flame extinguishes and
pop sound should be heard)
- Reaction gives out heat energy and catches flame (after K, K has a pretty lilac flame)
- Resulting solution turns Universal Indicator from green to purple, showing that it is alkaline.