Home Physics CS Accounting
Night More

Revise Now

Cambridge AS levels

Physics Computer Science Reference Questions

Programming and coding

Learn HTML Learn SEO


What do you want
to do?

Contact us Donate Support More Services Send a tweet Social media

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

trend of ionisation energy in period 3

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

ionisation trend down the group
The peaks represent the group 8 elements and it decreases as you go

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

The electronegativity of elements across a period

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 trend across a period
    This concept is very important as the elements on the left are the largest ones

    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

    ionic radius trend across a period
    This shows the trend for 2 periods. This shows how the pattern is the same for the two periods. This is the idea of periodicity

    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

    difference between the 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
    All of these produces a white solid oxide

    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.

    O2- + 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
    Chlorine is more electronegative than the other elements in its period

    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 boiling of elements across a period

    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

    The electric conductivity of elements across a period

    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

    End of Chapter Questions
    Collection of Topic Questions to understand the topic more.

    Try out

    Revisezone.com Copyright 2020