Periodic Trends (DP IB Chemistry): Revision Note

Periodicity

• The periodic table allows us to predict how certain properties change:

  • Across a period (left to right)

  • Down a group (top to bottom)

• These predictable patterns are known as periodic trends

Atomic radius

• The atomic radius is the distance from the nucleus to the outermost electron shell of an atom

  • It can be quite hard to determine exactly where the boundary of an atom lies, so a variety of approaches are taken such as half the mean distance between two adjacent atoms

  • This will vary depending on the type of structure and bonding, but it gives a comparative value for atoms

Atomic radius trend across a period:

• Atomic radius generally decreases across a period

• As you move across a period, the atomic number increases

  • This increases the number of protons in the nucleus

  • This results in an increased (positive) nuclear charge

• The extra electrons are added to the same principal quantum shell (energy level)

• The shielding between the nucleus and outer electrons is roughly constant

• Overall:

  • Increasing nuclear charge causes a stronger attraction between the nucleus and outer electrons

  • The stronger attraction means that the outer electrons are pulled closer to the nucleus

  • This results in the atomic radius decreasing

Atomic radius trend down a group:

• Atomic radius generally increases down a group

• As you move down a group, the number of principal quantum shells (energy levels) increases

  • This means that the outer electrons are further from the nucleus

  • This results in a weaker attraction between the nucleus and outer electrons

• Although nuclear charge increases, the electrons in the inner shells repel the outer electrons

  • This results in increased shielding between the nucleus and outer electrons

• Overall:

  • Increasing principal quantum shells causes the outer electrons to be further from the nucleus

  • There is increased shielding

  • So, the attraction between the nucleus and outer electrons decreases

  • This results in the atomic radius increasing

Diagram to show the trends in atomic radii

• The diagram shows that the atomic radius increases sharply between the noble gas at the end of a period and the alkali metal at the start of the next period

  • This is because the alkali metals in the next period have one extra principal quantum shell

  • This increases the shielding of the outermost electrons and therefore increases the atomic radius

Ionic radius

• The ionic radius is the distance from the nucleus to the outermost electron shell of an ion

Cations:

• The ionic radius generally decreases compared to the parent atom

• This is because electrons are lost, often resulting in the loss of an outer shell

• This leads to fewer electron–electron repulsions

• The remaining electrons are pulled in more strongly by the nucleus

• Overall:

  • Cations are smaller than their parent atoms because increased effective nuclear attraction pulls electrons closer

Anions:

• Ionic radius generally increases compared to the parent atom

• Extra electrons are gained, increasing electron–electron repulsion

• These repulsions outweigh the increased nuclear charge

• Overall:

  • Anions are larger than their parent atoms because the extra repulsion spreads the electron cloud out further

Ionic radius trend across a period:

• For cations:

  • Nuclear charge increases

  • Shielding stays roughly constant

  • So, the ionic radius decreases across a period

• After the point where anions form:

  • Ionic radius increases sharply due to added electron repulsion

  • Then decreases again as nuclear charge increases

Ionic radius trend down a group:

• Ionic radius increases

• As you move down a group, more electron shells are added

• Shielding increases and the outer electrons are further from the nucleus

• This leads to weaker nuclear attraction and larger ionic size

Diagram to show the trends in ionic radii

Ionisation energy

• The ionisation energy (IE) of an element is the amount of energy required to remove one mole of electrons from one mole of atoms of an element in the gaseous state to form one mole of gaseous ions

• Ionisation energies are measured under standard conditions which are 298 K and 100 kPa

• The units of IE are kilojoules per mole (kJ mol^-1)

• E.g. the first ionisation energy of calcium:

  • The first ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms

Ca(g) → Ca⁺ (g) + e^-            1st   ∆H IE = +590 kJ mol^-1

Ionisation energy trend across a period

• Ionisation energy generally increases across a period

• As you move across a period, the number of protons increases

  • This increases nuclear charge

• Electrons are added to the same principal energy level

  • Shielding remains roughly constant

• So, the outer electrons experience stronger attraction to the nucleus

• Overall:

  • More energy is needed to overcome this stronger attraction

  • So, ionisation energy increases across a period

Graph to show the trend in ionisation energies from H to Na

• There is a rapid decrease in ionisation energy between the last element in one period and the first element in the next period caused by:

  • The increased distance between the nucleus and the outer electrons

  • The increased shielding by inner electrons

  • These two factors outweigh the increased nuclear charge

Ionisation energy trend down a group:

• Ionisation energy generally decreases down a group

• As you move down a group, the number of principal energy levels increases

• The outer electrons are further from the nucleus and experience increased shielding

• The attraction between the nucleus and outermost electron decreases

• Overall:

  • Less energy is required to remove the outermost electron

  • So, ionisation energy decreases down a group

Ionisation Energy Trends across a Period & going down a Group Table

Electron affinity

• Electron affinity (EA) can be thought of as the opposite process of ionisation energy

• It is the amount of energy released when one mole of electrons is gained by one mole of atoms of an element in the gaseous state to form one mole of gaseous ions

• Electron affinities are measured under standard conditions which are 298 K and 100 kPa

• The units of EA are kilojoules per mole (kJ mol^-1)

First electron affinity

• The first electron affinity is usually exothermic for most nonmetals, as energy is released when an atom gains an electron

• Example of an exothermic first electron affinity: 

Cl (g) + e^– → Cl^– (g)                 ∆H = - 349 kJ mol^-1 

• However, some elements (e.g. Group 2, Group 12, and Group 15 elements) have positive first electron affinities, meaning the process is endothermic

  • This is due to the extra energy required to add an electron to a stable or half-filled subshell, where electron–electron repulsion is significant.

• Example of an endothermic first electron affinity: 

N (g) + e^– → N^– (g)                 ∆H = + 7 kJ mol^-1 

Second electron affinity

• However, the second electron affinity can be an endothermic process, e.g.

O^– (g) + e^– → O^2– (g)              ∆H = + 753 kJ mol^-1

• This is due to the fact that you are overcoming repulsion between the electron and a negative ion, so energy is required making the process endothermic overall

Electron affinity trends across a period:

• Electron affintity becomes more negative / exothermic across a period

• As you move across a period:

  • Nuclear charge increases

  • Atomic radius decreases

• So, the added electrons experience stronger attraction to the nucleus

• Overall:

  • Atoms more readily accept electrons

  • The energy released when an electron is added becomes more negative

Graph to show the electron affinities across a period

Electron affinity trends down a group:

• Electron affintity becomes less negative / exothermic down a group

• As you move down a group:

  • The number of principal quantum shells increases

  • Atomic radius increases

  • Shielding increases

• So, the added electrons are further from the nucleus and held less tightly

• Overall:

  • Less energy is released when an electron is added

  • So, electron affinity becomes less negative

• Note: Group 2 and group 15 elements show irregularities due to sublevel configurations and added electron repulsion

Graph to show the electron affinities down a group

Electronegativity

• Electronegativity is the ability of an atom to attract a pair of electrons towards itself in a covalent bond

• This phenomenon arises from the positive nucleus’s ability to attract the negatively charged electrons, in the outer shells, towards itself

• Electronegativity varies across periods and down the groups of the periodic table

Electronegativity across a period:

• Electronegativity increases across a period

• As you move across a period:

  • Nuclear charge increases with the addition of protons to the nucleus

  • Shielding remains the same across the period as no new shells are being added to the atoms

  • Atomic radius decreases

• This means that the bonding electrons are closer to the nucleus and feel a stronger attraction

• Overall:

  • Atoms more strongly attract bonding electrons

  • Electronegativity increases

Diagram to show the trend in electronegativity across a period

Electronegativity down a group:

• Electronegativity decreases down the group

• As you move down the group:

  • The nuclear charge increases as more protons are added to the nucleus

  • However, each element has an extra filled electron shell, which increases the shielding

  • The addition of the extra shells increases the distance between the nucleus and the outer electrons resulting in larger atomic radii

• Overall:

  • The effective nuclear attraction is weaker

  • So, atoms are less able to attract bonding electrons

  • Electronegativity decreases

Diagram to show the trend in electronegativities down a group