topic 3 - atomic radius and electronegativity Flashcards
Atomic Radius
The radius of an atom can change according to what is around it
The electrons are in constant motion so the overall shape and size of the atom can change
The only way to measure atomic radius is to measure the distance between two nuclei and divide by two
Trends across a period
As you go across a period, another proton is added each time
The additional electrons are being added to the same shell so doesn’t increase the radius
The extra proton in the nucleus causes an increase in the nuclear charge
This creates a stronger attraction to the surrounding electrons so pulls them closer
Therefore, atomic radius decreases across a period
Trends down a group
As you go down a group, atomic number is increasing
The extra electrons are being added to new shells
This means that as you go down a group, the atomic radius increases
Transition metals
the radii across the d-block are very similar as the increased nuclear charge is balanced by the extra shielding of the d-block electrons
Ionic Radius – trends down groups
Ionic radius increases as you go down a group as more shells are being added for each period
Ionic Radius – trends across periods
Cations (positive ions) are smaller than the corresponding atoms
They have lost the outer shell electrons so are smaller than the atom
They are isoelectronic – they all have the same electronic structure, i.e. Na+, Mg2+, Al3+
They have different numbers of protons though!
This means the pull from the nuclear charge is stronger across the period so the ionic radius decreases
Ionic Radius – trends across periods
Anions (negative ions) are larger than the corresponding atom
This is because the extra electrons push each other away (same charges repel)
Anions are also isoelectronic but there is still an increase in nuclear charge across the period
This means there is a stronger pull between the nucleus and outer electrons so the ionic radius decreases across the period
Electronegativity
This measures the tendency of an atom to attract a bonding pair of electrons
It increases as you go across a period
It decreases as you go down a group
Fluorine is the most electronegative element
Group 0 gases do not have electronegativity that can be reliably measured because they don’t make bonds
What causes electronegativity?
Electronegativity depends on:
Number of protons in the nucleus
Distance from nucleus to bonding electrons
Amount of shielding from inner electrons
what is shielding?
The shielding effect is a reduction in the attractive force between an electron and the nucleus of the atom caused by the inner shell electrons
The outer electrons are attracted by the nucleus, but also repelled by the electrons in the inner shells, making the overall attraction less
The more inner electron shells there are, the more the shielding effect so the less attraction there is between the outer electrons and the nucleus
Electronegativity – across the period
The bonding electrons are shielded by the same number of inner electrons
BUT proton number is increasing
This means electronegativity increases across the period
The Group 7 element will be more electronegative than the group 1 element
Electronegativity – down the group
As you go down a group, the atoms have more electron shells
This means there is more shielding and a greater distance between the nucleus and the bonding electrons
There is less pull between the nucleus and the bonding electrons
This means that electronegativity will decrease down the group
Electronegativity – other trends
There are some diagonal pairs of elements which have identical electronegativity
E.g. Beryllium and Aluminium
There is an increase in electronegativity from group 2 to group 3, but a decrease down the groups
The differences in electronegativity balance out so the two elements have the same electronegativity
This means they can react in a similar way and form similar bonds
Other pairs can do the same e.g. Lithium and magnesium
