What are the trends for atomic radius going across and down the periodic table?

In general, for the main group elements of the periodic table:

• Atomic radius1 increases down a Group, from top to bottom, of the Periodic Table.
• Atomic radius decreases across a Period, from left to right, of the Periodic Table2.

Note trends for ionic radii:

• ionic radius of a cation is less than atomic radius of the atom
• ionic radius of an anion is greater than atomic radius of the atom
• ionic radii of group 1 cations increases down the group
• ionic radii of group 17 anions increases down the group

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Trends in Atomic Radius in Groups of the Periodic Table

As you go down a Group in the Periodic Table from top to bottom, the number of energy levels or electron shells increases so the atomic radius of the elements increases.

Compare the number of occupied energy levels (electron shells) and the radius of the atom of elements in Group 1 and in Group 17 as given in each section below:

(A) Trends in the Atomic Radius of Group 1 (IA, Alkali Metals) Elements

Use the data in the table below for Group 1 elements to look for a pattern (or trend) in

• number of occupied energy levels (electron shells)
• atomic radius of the elements

As you go down Group 1:

• the number of occupied energy levels (or shells of electrons) increases from 2 to 6
• the radius of an atom of the element increases from 134 pm to 225 pm

Atomic radius increases as you go down the Group 1 elements from top to bottom as an additional energy level (electron shell) is being added to each successive element.

(B) Trends in the Atomic Radius of Group 17 (VIIA, Halogens) Elements

Use the data in the table below for Group 17 elements to look for a pattern (or trend) in

• number of occupied energy levels (electron shells)
• atomic radius of the elements

As you go down Group 17:

• the number of occupied energy levels (or shells of electrons) increases from 2 to 5
• the radius of an atom of the element increases from 71 pm to 133 pm

Atomic radius increases as you go down the Group 17 elements from top to bottom as an additional energy level (electron shell) is being added to each successive element.

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In general, the atomic radius of elements decreases as you go across a Period from left to right.

As we go across a Period from left to right, electrons are being added to the same energy level, the valence shell.
The increased nuclear charge, due to the positively charged protons, attracts all the negatively charged electrons more strongly so all the electrons are drawn in closer to the nucleus, in other words, the radius of each successive atom gets smaller across a period of the Periodic Table.

In the sections below we will consider the trends in the atomic radius of period 2 elements, and, of period 3 elements.

(A) Trends in the Atomic Radius of Elements in Period 2

Use the data given in the table below to find patterns (trends) in

• number of occupied energy levels (electron shells)
• charge on the nucleus (nuclear charge)
• atomic radius of the elements

Can you see these patterns (trends):

• number of occupied energy levels (electron shells) remains the same (2 occupied electron shells) across the period
• charge on the nucleus (nuclear charge) increases from +3 to +10 across the period from left to right
• atomic radius of the elements decreases from 134 pm to 69 pm across the period from left to right

Atomic radius generally decreases across Period 2 from left to right as the nuclear charge increases.

(B) Trends in the Atomic Radius of Elements in Period 3

Use the data given in the table below to find patterns (trends) in

• number of occupied energy levels (electron shells)
• charge on the nucleus (nuclear charge)
• atomic radius of the elements

Can you see these patterns (trends):

• number of occupied energy levels (electron shells) remains the same (3 occupied electron shells) across the period
• charge on the nucleus (nuclear charge) increases from +11 to +18 across the period from left to right
• atomic radius of the elements decreases from 154 pm to 97 pm across the period from left to right

Atomic radius generally decreases across Period 3 from left to right as the nuclear charge increases.

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• Cations are smaller than their respective atoms as electrons are removed from the highest energy level (valence shell) while the positive nuclear charge remains the same thereby increasing the attraction between the remaining electrons and the nucleus resulting in a reduction in the size of the radius of the cation.
• Anions are larger than their respective atoms as electrons are added to the highest energy level (valence shell) the repulsion between the negatively charged electrons increases the ionic radius.
• Down a Group of the Periodic Table from top to bottom, comparing ions with a similar charge, the ionic radius increases as the number of completed energy levels increases.

Let's compare the radius of some elements' atoms with their respective ions. Group 1 elements form cations with a charge of +1

Group 17 elements form anions with a charge of -1

(A) Comparison of Atomic and Ionic Radius of Group 1 (IA, alkali metals) Elements

Atoms of group 1 elements (M) lose 1 electron (e-) from their valence shell of electrons to form a cation with a charge of +1 (M+)
We can wrote a chemical equation to represent this for any group 1 element (M) as shown below

We can write a specific equation to represent the formation of each cation of each group 1 element as shown below:

Consider the data in the table below. Can you see any patterns (or trends) in the data?

Compare each atom and its respective cation with regards to:

• atomic and cationic radius
• number of occupied energy levels (electron shells)

Did you notice that:

• the cationic radius of an element is always less than its atomic radius?

  (a) ionic radius Li+ (74) < atomic radius Li (134)
  (b) ionic radius Na+ (102) < atomic radius Na (154)
  (c) ionic radius K+ (138) < atomic radius K (196)
  (d) ionic radius Rb+ (149) < atomic radius Rb (211)
  (e) ionic radius Cs+ (170) < atomic radius Cs (225)

• the cation has 1 less occupied energy levels (electron shells) ?

  (a) Li has 2 electron shells occupied, Li+ has 2 - 1 = 1 electron shell occupied
  (b) Na has 3 electron shells occupied, Na+ has 3 - 1 = 2 electron shells occupied
  (c) K has 4 electron shells occupied, K+ has 4 - 1 = 3 electron shells occupied
  (d) Rb has 5 electron shells occupied, Rb+ has 5 - 1 = 4 electron shells occupied
  (e) Cs has 6 electron shells occupied, Li+ has 6 - 1 = 5 electron shells occupied

As you go down the group from top to bottom, the ionic radius of each cation increases, just as it does for the radius of each atom because an energy level (electron shell) is being added to each successive atom (or ion).

However, if an electron is removed from a Group 1 atom it is removed from the highest energy level (valence shell), so this effectively reduces the number of occupied energy levels (or electron shells).
The same nuclear positive charge is now acting on fewer negatively charged electrons so these electrons are drawn in closer towards to the nucleus and the radius of the cation will be less than the radius of the atom.

(B) Comparison of Atomic and Ionic Radius of Group 17 (VIIA, halogen) Elements

1 electron (e-) can be added to the valence shell of the atoms of group 17 elements to form an anion with a charge of -1 (X-)
We can write a chemical equation to represent this for any group 17 element (X) as shown below

We can write a specific equation to represent the formation of each anion of each group 17 element as shown below:

Consider the data in the table below. Can you see any patterns (or trends) in the data?

Compare each atom and its respective anion with regards to:

• atomic and anionic radius
• number of occupied energy levels (electron shells)

Did you notice that:

• the radius of the anion is always greater than the radius of the atom?

  (a) radius of F- (131) > radius of F (71)
  (b) radius of Cl- (181) > radius of Cl (99)
  (c) radius of Br- (196) > radius of Br (114)
  (d) radius of I- (220) > radius of I (133)

• each electron is being added to the same energy level (electron shell)?

  (a) number of occupied electron shells F = number of occupied electron shells F- = 2
  (b) number of occupied electron shells Cl = number of occupied electron shells Cl- = 3
  (c) number of occupied electron shells Br = number of occupied electron shells Br- = 4
  (d) number of occupied electron shells I = number of occupied electron shells I- = 5

Ionic radius increases as you go down the Group 17 elements from top to bottom as an additional energy level (electron shell) is being added to each successive element, just as it is for the neutral atom.

But what causes the radius to expand when an electron is added to the valence shell of the original atom?

The charge on the nucleus of the atom caused by the number of positively charged protons in the nucleus does not change when an electron is added to an electron shell. But the number of negatively charged electrons in that electron shell does change! And it means that there is a greater repulsion between all the electrons.

The effect of this repulsion between electrons is to increase the radius of the anion compared to the neutral atom.

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Footnotes

1. The values for atomic radii provided here are approximations derived from interatomic-distance measurements.
There are a number of different ways of approaching the measurement of atomic and ionic radii. We are not going to discuss any of them, we are just going to use some nicely behaved approximations to illustrate some general trends.

2. We will be excluding transition metals, lanthanoids (lanthanides) and actinoids (actinides) from the disucssion

3. Common units for the reporting of atomic radii are the picometre (as used here), the nanometre, and the angstrom.
1 pm = 1 picometre = 10-12 metre
1 Å = 1 angstrom = 10-10 metre
1 nm = 1 nanometre = 10-9 metre Conversion example for the atomic radius of lithium,

134 pm = 134 × 10-12 m = 0.134 × 10-9 m = 0.134 nm

4. Ionic radius based on a coordination number of 6