This page covers the following topics:
1. Atomic radius
2. Melting point
3. 1st ionisation energy
4. 2nd 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
Melting point is the temperature at which a substance changes its physical state from solid to liquid. There is no strict trend that all periods or groups follow in terms of melting point. The melting point of an element depends on the type of bonding it forms.
The melting point of metals increases with valency, as its bonding strength depends on the volume of delocalised electrons present.
Covalent bonds are the strongest and they are broken when substances with strong covalent structures melt. Carbon therefore has the highest melting point of all elements.
Molecular elements are only held together by weak van der Waals forces and no covalent bonds are broken when they melt. Thus, they have the lowest melting point. The larger the molecule, the more van der Waals interactions it can form. For larger molecules (P₄, S₈) stronger van dear Waals forces have to be broken (in comparison to Cl₂, Ar), requiring more energy during melting and resulting in a higher melting point.
You may use the following link to visualise the trends of melting point in the periodic table: https://ptable.com/#Properties/MeltingPoint
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
2nd ionisation energy is the amount of energy required to remove the second valence electron from a gaseous atom. 2nd ionisation can be presented using a half-equation that shows a singly charged ion X⁺ forming X²⁺ and an electron: X⁺(g) → X²⁺(g) + e⁻.
It follows a similar trend to 1st ionisation energy, depending on atomic radius and electron configuration. If an electron during the second ionisation is removed from an orbital with a lower energy, the 2nd ionisation energy is going to be higher. If an electron to be removed is paired, 2nd ionisation energy will be lower.
The 2nd ionisation energy is generally a lot bigger than the 1st ionisation energy for the same element, as there is a reduced electron repulsion. A singly charged element, has less electrons than the neutral element. Therefore, there is a weaker repulsion between electrons and they are held closer to the nucleus. This makes it harder to remove an electron from an initial ion, thus more energy is required.
You may use the following link to visualise the trends of the 2nd ionisation energy in the periodic table: https://ptable.com/#Properties/Ionization/2nd
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.
How does electronegativity change going towards the right across a period?
Explain why this sketch is an incorrect depiction of a bond between fluorine and boron.
Write a half-equation for the 2nd ionisation of gallium.
Which metal has the largest melting point: potassium, calcium or scandium?
Why is the 2nd ionisation energy of sodium (4562 kJ/mol) much higher that the 1st ionisation energy of sodium (498 kJ/mol)?
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