Periodicity Notes

Across any period in a periodic table, there is an recurring pattern or trend for each property. We will discuss the main trends and patterns you need to know. We will use period 3 as our main example because this is what most exams follow

Ionisation Energies across the period

Across the period, the Ionisation energies generally increases. This is because, the shielding effect is more or less the same as electrons are added to the same principle Quantum shell and the atomic radius slightly decreases. Also the nucleur charge or the proton number increases causing a greater nucleur charge and attraction forces between the nucleus and the valence electrons. So more energy is required

The above explanation is enough to score marks. However, this is the general trend and so there are some exception you need to know

Exceptions to the rule

The exception are only for group 2,3 and group 5,6 elements

As you guessed it, group 3 elements has a slightly lower ionisation energies than the group 2 elements. Why? We will see the answer below

The electrons of group 3 elements is added to the P subshell instead of the S orbital. So due to slight increase in atomic radius and shielding effect being more significant than the increase in nucleur charge. The effective nucleur charge decreases and so lower attractive force

If you are answering a question, ensure that you mention the principal quantum shells of the Subshells you are talking about. For example, 3P and 3S. Also remember to use the word "significant" in your answer as it is very important

Now we will go to the group 5, 6 exception

The group 6 has a lower ionisation energy compared to group 5 elements. The reason is a bit different and it includes the 4th factor

For group 6 elements, the electron is paired with another electron whereas for group 5 elements the p subshell is half filled with individual electrons. So due to the spin-pair repulsion of the electrons, it reduces the effective nucleur charge or attraction forces and so less energy is required

This idea is very important, that is why remember that group 5 elements have a half filled P subshell

This pattern works for almost across any period. So that's why it is part of the periodicity notes!

Trends down the group

To simply put it the ionisation energies decrease down the group. Why? We will see the answer

The increase in atomic radius and the increase in shielding effect due to increase in inner core electrons is more significant than the increase in nucleur charge or proton number. So the effective nucleur charge and the attraction forces between the electrons and the nucleus decreases and so the ionisation energy decreases

It is that simple and that's all you need to know! Make sure you use the word significant!

Electronegativity Trends

There is a clear increase in electronegativity across the period. Note that this does not follow the trend as the ionisation energy but the reasons are almost identical

Also NOF are the most electronegative elements

How is Atomic Radius determined

For different types of bonding the atomic radius is found differently:

Metallic bonding

Half of the distance between the centers of the two positive cations

Covalent bonding / Diatomic

Half of the distance between the centers of the two nuclei that is covalently bonded.

This will not give a precise radius as the orbitals are overlapping

Monoatomic

Half of the distance between the centers of two atoms touching

Trends in Atomic Radius

Atomic radius is a periodicity section but, we will still discuss it as it is related to ionisation energies

All you need to know that across a period, the atomic radius of elements decreases. Why? This is because, due to the increase in nucleur charge and shielding effect is more or less constant, there is greater attraction forces between the valence electrons and the nucleus and so greater pull. So the radius decreases

You will need to remember this graph and this is similar for any period

Exams question by asking to compare the atomic radius of 2 elements. For example, which one is larger in size, Na or Al. Even though it is tempting to think that because Al has more subshells, the atomic radius of Aluminium is less than Sodium. So Sodium is larger in size

Trends in Ionic Radius

When comparing Ionic Radius, it is a bit different. Usually when elements form ions, they form the most stable electronic configuration

The Ionic radius of elements across a period also decreases but the anions have a larger radius than the cations. Remember the graph!

This is because, the anion has an extra principle quantum shell to hold more electrons and also there is greater shielding effect from inner core electrons

But for both Cations and Anions the Ionic radius decreases because the nucleur charge increases and the shielding effect is constant and so greater attraction and pull

Comparison between Ionic and Atomic Radius

When we compare the Ionic or the atomic radius of a particular element, you need to consider whether it is a Anion or a cation!

Simply the cations has a smaller radius than the corresponding atom. Why? This is because, the cation has less electrons than the atom and so there is less repulsion force between electrons causing a greater EFFECTIVE nucleur charge on the valence electrons

The Anions has a larger radius than the corresponding atoms because, the anions have more electrons and so greater repulsion forces. This leads to a lower EFFECTIVE Nucleur charge on valence electrons so the radius is more larger

That's all you need to know. Please use the explanations above as they are from the markshemes!

Reactions with Oxygens

We will see how elements across the period react with oxygen. We will take period 3 as the example

Na - Burns with a yellow flame
Mg - Burns with a bright white flame
Al - Slowly oxidises
Si - Slowly oxidises
P  - Burns with a yellow flame
S - Burns with a blue flame

Solubility of oxides

We will now see how the solubility of oxides changes across a period

Compound	Solubility
Na2O	Very Soluble
MgO	Slightly Soluble
Al2O3	Insoluble
SiO2	Insoluble
P2O5	Soluble
SO2	Soluble

pH of Oxides

The pH of the oxide depends on whether the compound is ionic or covalent

Compound	Nature
Na2O	Basic
MgO	Basic
Al2O3	Ampthoreic
SiO2	Acidic
P2O5	Acidic
SO2	Acidic

You need to know that Aluminium oxide and Silicon dioxide does not dissolve but does react

When a metal oxide such as sodium dissolve in water, OH^- ions are released.

O^2- + H₂O → 2OH^-

That is why Na₂O and MgO are basic...

Period 3 Chlorides

There is a clear trend in the acidity of chlorides in water across the period

Compound	pH
NaCl	7
MgCl2	6
Al2Cl6	4
SiCl4	3
PCl5	3
SCl2	3

What you need to remember is that all Chlorides except the group 1 chlorides are acidic. Group 1 Chlorides are always neutral, the acidicity increases across the period and you need to know these pH values for each element

Another Point to remember is that, the chlorides are HYDROLYSED by water which is a reaction that makes the solution acidic. Only Sodium Chloride does not hydrolyse but just dissolves in water

Oxidation Number of Chlorides & Oxides

We need to remember that the more electronegative atoms are always negatively charged. We will see the Oxidation of each element in their chlorides:

Compound	Oxidation Number
NaCl	+1
MgCl2	+2
Al2Cl6	+3
SiCl4	+4
PCl5	+5
SCl6	+6

We will see the oxidation number of Oxides. It is very similar

Compound	Oxidation Number
Na2O	+1
MgO	+2
Al2O3	+3
SiO2	+4
P2O5	+5
S02	+4

Trends in Bonding of Elements

The physical properties usually depend on the bonding involved in each group. Usually, there is a same type of bonding for each group

Group	Bonding
1	Metallic
2	Metallic
3	Metallic
4	Giant Covalent
5	Simple molecular
6	Simple molecular
7	Simple molecular
8	Monoatomic

By using this, we can predict the trend of the melting and electric conductivity across the period.

Melting and Boiling points

The melting point from Group 1 to Group 5 increases. We will see why

The Group 1 , 2 & 3 have metallic bonding but from group 1 to 3, the charge of the ions formed by the metals increases ( Na⁺, Mg²⁺, Al³⁺ ). So higher the charge, the higher the charge density and greater electrostatic forces between metal ions. This makes it much more stronger and have a higher melting point.

For group 4, as it is a giant covalent structure (full of covalent bonds), this makes the structure extremely strong and have the highest melting & boling point

For group after the Group 4, the melting point decreases as they are simple molecular strucures. However, the group 6 element has a slightly higher boiling and melting point than its usual trend.

Electric Conductivity

This also depends on the bonding. From group 1 to 3, the electric conductivity increases as the metal ions have a greater charge and so there is more free delocalised electrons to carry charge.

From group 3 onwards the electric conductivity decreases. This is because giant covalent and simple molecular structures do not have free mobile electrons to carry charge