O Level Chemistry Periodic Table Notes (Revision)

You are about to go through O Level Chemistry periodic table notes. This is a very important topic as plenty of questions are appearing recently in the CIE.

So, let’s dive straight in the topic without further introductions.

What is a Periodic Table?

A periodic table is a listing of elements that are regulated in order of increasing atomic (proton) number.

The elements in the periodic table are divided into periods and groups. How do we define a period and a group in a periodic table? Let’s find out:

Groups:

• It is a vertical column of elements.
• They run from top to bottom and there are eight groups in total (in the periodic table).
• Group V||| is also referred as group 0.

Periods:

• Periods are horizontal rows.
• They run from left to right and there are 7 periods in total (1 to 7).

The Periodic Trends

The elements in the periodic table can be classified based on their metallic and non-metallic properties. You should also know that there is an imaginary diagonal line in the periodic table.

The elements located next to this line are metalloids, which means that they have the properties of both metals and non-metals.

Metals can be found at the left side of the periodic table while the non-metals are found towards the right of the periodic table. Therefore, from left to right in a periodic table, there is a decrease of metallic properties and increase of non-metallic properties.

Note: Remember that the metals are basic in nature and the non-metals are acidic in nature. What about the metalloids? Well, they are amphoteric in nature.

The change of metallic properties down a group:

If you move down a group (vertically), you will notice an increase in the metallic properties and a decrease in the non-metallic properties.

This is because the valence (outermost) electrons move far away from the nucleus and therefore, the tendency to lose electrons increases. The easily the metal loses electrons, the more reactive it is.

Electronic Structure:

The proton number helps us to find out the electronic structure of an element. Moreover, the atomic number of an element also allows to to find out the group and period number of an element.

Let’s understand this with the help of an example:

Let’s use sodium as an example. Sodium has a proton number of 11. Since the number of protons is equal (=) to the number of electrons, therefore, sodium has 11 electrons.

Thus, the electronic configuration of an sodium atom is:

2, 8, 1

• From this information, we can deduce that the period number of sodium is 3, why?

The number of electron shells is equal to the period number of an element.

This means that an element with period number 3 will have 3 electron shells!

• We can deduce that the group number of sodium is 1, how?

Look, the number of electrons in the outermost shell is similar to the group number of an element. Therefore, sodium has 1 valence electron which means its group number is 1.

The number of valence (outermost) electrons is equal to the group number of the element.

Note: The elements with same valence electrons have similar chemical properties. We can also say that the elements in the same group share similar chemical properties.

Further reading:

• Deformation Notes O Level Physics
• Best Dynamics Notes (Forces) | GCE O Level Physics (5054)
• O Level Report Writing Sample (Directed Writing)

Now let’s discuss some important groups in the periodic table individually:

Group 1 – Alkali Metals:

The elements in the group one of the periodic tables are known as the alkali metals. The elements in group one are:

1. Lithium
2. Sodium
3. Pottasium
4. Rubidium
5. Cesium (Cs)
6. Francium (Fr)

Physical Properties of Alkali Metals:

• Their melting and boiling point is low and it decreases down the group, why?

This is because the metallic bonds are easily disrupted by heat as they are orderly arranged.

This orderly arrangement can be easily disturbed by heat which explains why Alkali metals have low melting and boiling point.

• They can be easily cut because they are soft. The reason being is their orderly arrangement which can be easily disrupted if force is applied!

• They have low densities which allow them to float on water. However, down the group, the density increases because as atomic radius increases, the volume increases (the number of valence electrons increase).

Alkali metals are very reactive!

Alkali Metals – Chemical Properties:

• Alkali metals get oxidised and therefore, they work as strong reducing agent (element or a compound that itself gets oxidised).

In other words, Alkali metals lose electrons and become positively charged ions.

• Alkali metals react with water (cold) to form hydrogen and an alkali.

Since going down the group increases the reactivity, the reaction of sodium and potassium with cold water can be violent because they are more reactive as compared to Lithium.

Halogens (Group V|| Elements)

Halogens are those elements that are located in the group seven of the periodic table. The group seven of the periodic table includes elements such as:

1. Chlorine
2. Bromine
3. Iodine
4. Astatine
5. Fluorine

Halogens are non-metals and they have low melting and boiling point. This is because the diatomic covalent bonds can easily be broken.

If you move from top to bottom in the group V|||, you will notice that the boiling point increases and the colour of the elements become darker.

Halogens are very reactive non-metals because they have seven electrons in their outermost shell. They only need one electron to complete their outermost shell. This makes them fairly recative.

They are known as Halogens because they produce salts when they react with Sodium Carbonate and the word “Halogen” is derived from the Greek word hal– which means salt and gen– refers to produce.

See: O Level Formal Letter Sample (Directed Writing)

Chemical Properties of Halogens:

• A more reactive halogen can replace a less reactive halogen in a displacement reaction.

For example; chlorine is more reactive than bromine as it is higher than bromine in the group V|| of the periodic table. Therefore, it will displace bromine in the below reaction.

Cl2 + 2KBr → 2KCl + Br2

• Halogens are strong oxidising agents because they themselves get reduced (being non-metals, they gain electrons).

Noble Gases: Group V||| Or Group 0:

Noble gases are also known as inert gases (because they are un-reactive). Some of the noble gases are helium, neon, argon, krypton, xenon and radon.

Noble gases are colourless at room temperature and they are monoatomic (stable as single atoms).

They are unreactive (because they have stable duplet or octet electronic configuration) and insoluble in water. The noble gases are present in the Earth’s atmosphere.

Why is it important to study about noble gases?

In O Level chemistry periodic table notes, you should know that the noble gases have many uses in our day to day life (based upon their chemical properties). Some of them are listed below.

Some uses of noble gases:

• Xenon is used in vehicle headlamps and lighthouses due to its very bright light. Similarly, it is also used in producing lights of different colours.

• Argon is used to provide an inert (chemically inactive) atmosphere. It is also used in transistors and electric light bulbs.

• The advertisement signs are made with neon gas.

• Helium is used in air tanks (for deep sea-diving), balloons and to cool superconducting magnets (electromagnet made from coils of superconducting wire).

Note: Noble gases conduct electricity at low pressure. They are non-flammable and do not have any colour, flavour or odour. Apart, they have a low melting point and electronegativities (tendency to form bonds because they are fairly unreactive).

Transition Metals:

The elements (manganese, iron, copper and chromium etc) located between group 2 and 3 (the middle portion of periodic table) are known as transition metals.

They have high melting and boiling point (high density also) as their metallic bonds are very strong and thus, high heat can overcome the forces of attraction.

Moreover, transition elements have variable oxidation state (valency). Well, what does that mean?

The ions of Iron are Fe²⁺ and Fe³⁺. Thus, Iron has two oxidation states. Since it has more than one oxidation state, it is said to be having variable oxidation states.

Same is the case with other transition metals such as manganese and copper.

Furthermore, transition metals form coloured compounds and their compounds are very good catalysts. E.g Iron (used in Haber process) and Nickel (used in the manufacture of margarine).

The reason why they are very good catalysts is their variable oxidation state that allow the industrial processes to be happen at a quicker pace.

Why transition metals form coloured compounds is not required by your syllabus because it is meant for higher studies. However, let me tell you that the d orbitals in them are half-filled or unfilled.

This is the “Crystal field theory” which refers to the movement of d orbitals which is the reason why transition metals form coloured compounds.

It is to note that transition metals are used in the making of many important alloys. For instance, stainless steel is an alloy of chromium and iron (resists corrosion very well). The summary of the properties of transition metal is:

• They have a very high melting and boiling point.

• They are good conductors of electricity and heat.

• They are hard and have usually high densities.

• They can be hammered into different shapes.

Conclusion:

With this, O Level Chemistry Periodic table notes have been delivered. If you want further guidance regarding the topic, do leave a comment.

Thank You very much for staying with me in the end. This topic is very important from MCQ’s point of view. So, remember to practice some past papers so that your revision becomes even more effective.

Stay tuned for more and some of the important points that have been covered in this topic are listed below for your reference.

1. Introduction to the periodic table
2. Periodic trends
3. Alkali metals
4. Electronic configuration
5. Halogens and noble gases
6. Transition metals