Periodic table trends for SQA Higher Chemistry
This page covers the following topics:
1. Covalent radius
3. Ionisation energy
As you move towards the right across a period, the atomic number and nuclear charge gets bigger. This causes a greater attraction between the nucleus and the electrons, pulling them closer to each other. Thus, the size of an atom and its radius gets smaller.
As you go down a group, the atomic radius of atoms gets bigger as they gain additional shells of electrons. The number of shells is more important than the position within a period while considering the size of an atom.
You may use the following link to visualise the trends of the atomic radius in the periodic table: https://ptable.com/#Properties/Radius/Calculated
The electronegativity of an element describes how readily it attracts a shared pair of electrons towards itself. For example, if an atom of the least electronegative element, francium, were to form a bond with an atom of the most electronegative element, fluorine, electrons that participate in bonding between them would be primarily surrounding the fluorine atom.
The electronegativity of an element increases as you move from the bottom left to the top right corner of the periodic table. There are many mathematical definitions of electronegativity, the most commonly used is the Pauling scale.
The 1st ionisation energy of an element is the amount of energy required to remove the first valence electron from a gaseous atom. 1st ionisation can be presented using a half-equation that shows a neutral atom X forming X⁺ and an electron: X(g) → X⁺(g) + e⁻.
As the atomic radius decreases, the nucleus has a tighter hold on its electrons, so more energy is required to remove them. Therefore, the 1st ionisation energy increases going towards the right across a period and decreases going down a group.
There are exceptions to the general trend. Electrons present in a sub-shell further away from the nucleus or at a higher energy would require less energy to be removed from an atom. Paired electrons are easier to remove from an atom than an unpaired ones from the same sub-level. Reduced energy required results in a lower 1st ionisation energy in both cases.
You may use the following link to visualise the trends of the 1st ionisation energy in the periodic table: https://ptable.com/#Properties/Ionization/1st
How does electronegativity change going down a group of the periodic table?
Electronegativity decreases going down a group of the periodic table.
Why does the radius of sodium atom decrease when it loses an electron and becomes Na⁺ ion?
The one electron that sodium atom loses to become Na⁺ ion is the only electron in its atomic outer shell. Losing this electron means losing its outer shell. We know that as you go down a group in the periodic table, the atomic radius of elements gets bigger as they gain additional shells of electrons. Thus, sodium atom losing an electron shell results in a reduction of its radius.
loss of an electron shell → reduction in the radius of the particle
Explain why this sketch is an incorrect depiction of a bond between fluorine and boron.
Fluorine is more electronegative than boron. The shared electrons should be held closer to fluorine.
Due to higher electronegativity of fluorine, electrons would be closer to it.
Write a half-equation for the 1st ionisation of silver.
In the first ionisation half-equation silver would go from a neutral gaseous atom to a gaseous Ag⁺ and would produce an electron.
Ag(g) → Ag⁺(g) + e⁻
Explain the general trend shown in the diagram.
The general trend from going towards the right in the third period shows an increase in 1st ionisation energy. This can be explained by the decrease in atomic radius as you go along the period. The nucleus has a tighter hold on the valence electrons, so more energy is needed to remove them.
decrease in atomic radius towards the right → stronger attraction between valence electrons and a nucleus → higher ionisation energy
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