This is one of the shortest topics in this module. The other being Group VII elements.
Let's get started.
At the end of this topic, you will know the following information about Group II elements:
- Trends in properties
- Redox reactions
- Reactions with oxygen, water and hydrochloric acid
- Thermal decomposition of the carbonates
- Uses of Group II compounds
Trends in properties Edit
The Group II elements are powerful reducing agents.
A reducing agent is the compound that gets oxidised in the reaction and, therefore, loses electrons.
M = Mg, Ca, Sr,Ba --> I will be using 'M' as the general symbol for a Group II element in this topic.
As I said earlier, they are powerful reducing agents. M(s) ----> M2+(aq) + 2e-
A reducing agent 'loses electrons'.
Another term for Group II elements are Alkaline Earth Metals
All Group II elements have 2 electrons in their outer shell. They generally lose these two outershell electrons in order to react and, by doing so, they form M2+ ions.
These are the properties you will have to know for your exam
As you go down the Group:
- Atomic radius increases
- First ionisation energy decreases
- Chemical reactivity increases
- Electronegativity decreases
You need to be able to explain these properties in the exam.
- Why does atomic radius increase?
- This happens because as you go down the Group, the elements have more electrons and therefore, need more orbitals in which to store these electrons. These orbitals get filled up gradually by the usual method: "1s2 2s2 2p6 3s2 3p6 4s2 3d10 etc etc...
- These orbitals surround the nucleus. As you go down the Group, you get more orbitals surrounding the nucleus (remember, all the orbitals are 'stacked' on top of each other) and therefore, the size of the atom increases.
- Also, the further an orbital gets from the nucleus, the even more further it gets because of the sudden decrease in nuclear attraction from the nucleus.
- Why does first ionisation energy decrease?
- This happens because as you go down the Group, the number of orbitals increases and they get further away from the nucleus. This decreases the nuclear attraction between the orbital and the nucleus so it makes it easier for another element to remove an electron from the outer orbital as the outer orbital is hardly attracted to the nucleus.
- Also, another reason the first ionisation energy decreases is because, of the shielding effect. All orbitals are filled up with electrons and these electrons are constantly repelling each other. The further out an orbital is from the nucleus, the greater the shielding effect on it and the lower the nuclear charge. Therefore, it's easier to remove an electron.
- Exam tip
- Whenever you are asked a question about explaining ionisation energies, always include the following 3 points.
- Distance from nucleus
- Shielding effect
- Proton:Electron Ratio
- Why does chemical reactivity increase?
- This happens because the ionisation energy decreases. All chemical reactions take place due to the transfer of electrons. Ionisation energy is the energy required to remove one electron from each atom in a one mole sample of gaseous atoms to form one mole of gaseous unipositive atoms.
Therefore, if its easier to remove an electron, then it should mean the it will be easier to react.
- Why does electronegativity decrease?
- The ability of M to attract the electrons in a bonding pair decreases as you go down the group because the bonding orbitals are further from the nuclear charge and shielded by the inner shells. Note that group II metals form mostly ionic compounds because the electronegativities are significantly lower than elements such as oxygen and chlorine. Beryllium has the highest electronegativity in Group II and, as you might predict, it forms the chloride with most covalent character.
Redox reactions Edit
A Redox reaction is a reaction in which Oxidation and Reduction take place at the same time. Originally Oxdation was defined as the gaining of oxygen by an element and Reduction as the loss of oxygen.However, now Oxidation is defined as the loss of Electrons and Reduction as the gain of electrons.
Oil Rig Edit
A simple way of remembering oxidation and reduction is OILRIG
Oxidation is loss of electrons
Reduction is gain of electrons
Reactions with oxygen, water and hydrochloric acid Edit
A reaction can only be regarded as a redox reaction if oxidation and reduction occurs simultaneously<
Thermal decomposition of the carbonates Edit
Carbonates of group II metals decompose on heating to give carbon dioxide.
MCO3 -----------------> MO + CO2
This is used to convert chalk (calcium carbonate) to quicklime, calcium oxide. Quicklime can be used to make a simple mortar for building so historically lime kilns and small quarries were to be found all over England. When roasted with clay quicklime is used to make cement.
The CO2 produced by this reaction is also used to manufacture Sodium Carbonate (see Solway process).
Quicklime reacts vigorously with water to form Calcium Hydroxide (slaked lime). This is a very exothermic reaction.
CaO + H2O ---------------> Ca(OH)2
Slaked lime can purchased in garden centres as a soil conditioner, farmers spread it on the land to reduce acidity.
Uses of Group II compounds Edit
Thermal decomposition of the carbonates The carbonates of the group two metals decompose when heated to produce CO2. This is then used to change chalk to calcium oxide, which is also known as quicklime. The reaction to create the CO2 is:
MCO3 -----------------> MO + CO2
This calcium oxide is used in construction for very simple mortar. If it is roasted with clay, it creates cement.
The CO2 produced by this reaction is also used to manufacture Sodium Carbonate (see Solway process). Quicklime reacts vigrously with water to form Calcium Hydroxide (slaked lime). This is a very exothermic reaction. CaO + H2O ---------------> Ca(OH)2 Slaked lime can purchased in garden centers as a soil conditioner, farmers spread it on the land to reduce acidity.
Summary The Group II elements are powerful reducing agents. A reducing agent 'loses electrons'. Another term for Group II elements are Alkaline Earth Metals All Group II elements have 2 electrons in their outer shell.