Periodic Trends (SL IB Chemistry)

Atomic radius

• The atomic radius of an element is a measure of the size of an atom
• It is the distance between the nucleus of an atom and the outermost electron shell
• 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 diagram

The atomic radius of an atom is the typical distance between the nucleus and the outermost electron shell

Trends in atomic radii

• Atomic radii show predictable patterns across the periodic table
  • They generally decrease across each period
  • They generally increase down each group
• These trends can be explained by the electron shell theory
  • Atomic radii decrease as you move across a period as the atomic number increases (increased positive nuclear charge) but at the same time extra electrons are added to the same principal quantum shell
  • The larger the nuclear charge, the greater the pull of the nuclei on the electrons which results in smaller atoms
  • Atomic radii increase moving down a group as there is an increased number of shells going down the group
  • The electrons in the inner shells repel the electrons in the outermost shells, shielding them from the positive nuclear charge
  • This weakens the pull of the nuclei on the electrons resulting in larger atoms

Diagram to show the trends in atomic radii

Trends in the atomic radii across a period and down a group

• The diagram shows that the atomic radius increases sharply between the noble gas at the end of each period and the alkali metal at the beginning of the next period
• This is because the alkali metals at the beginning of 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 of an element is a measure of the size of an ion
  • The trend down a group is the same as atomic radius - it increases as the number of shells increases
  • The trend across a period is not so straightforward as it depends on whether it is positive or negative ions are considered
  • Ionic radii increase with increasing negative charge
  • Ionic radii decrease with increasing positive charge
• These trends can also be explained by the electron shell theory
  • Ions with negative charges are formed by atoms accepting extra electrons while the nuclear charge remains the same
  • The extra electrons experience repulsion with the other valence electrons which increases the ionic radius
  • The greater the negative charge, the larger the ionic radius
  • Positively charged ions are formed by atoms losing electrons
  • The nuclear charge remains the same but there are now fewer electrons which undergo a greater electrostatic force of attraction towards the nucleus which decreases the ionic radius
  • The greater the positive charge, the smaller the ionic radius

Diagram to show the trends in ionic radii

Trends in the ionic radii across a period and down a group

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

Trends in ionisation energy

• Ionisation energies show periodicity
• As could be expected from their electronic configuration, the Group 1 metals show low IE whereas the noble gases have very high IEs
• The first ionisation energy increases across a period and decreases down a group and is caused by four factors that influence the ionisation energy:
  • Size of the nuclear charge: the nuclear charge increases with increasing atomic number, which means that there are greater attractive forces between the nucleus and outer electrons, so more energy is required to overcome these attractive forces when removing an electron
  • Distance of outer electrons from the nucleus: electrons in shells that are further away from the nucleus are less attracted to the nucleus so the further the outer electron shell is from the nucleus, the lower the ionisation energy
  • Shielding effect of inner electrons: the shielding effect is when the electrons in full inner shells repel electrons in outer shells preventing them to feel the full nuclear charge so the greater the shielding of outer electrons by inner electron shells, the lower the ionisation energy
  • Spin-pair repulsion: paired electrons in the same atomic orbital in a subshell repel each other more than electrons in different atomic orbitals; this makes it easier to remove an electron (which is why the first ionisation energy is always the lowest)

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

A graph showing the ionisation energies of the elements hydrogen to sodium

Ionisation energy across a period

• The ionisation energy across a period increases due to the following factors:
  • Across a period the nuclear charge increases
  • The distance between the nucleus and outer electron remains reasonably constant
  • The shielding by inner shell electrons remains the same

• 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 down a group

• Although going down a group the nuclear charge increases, the ionisation energy down a group decreases and it is due to the following factors:
  • The distance between the nucleus and the outer electron increases
  • The shielding by inner shell electrons increases
  • The effective nuclear charge is decreasing as shielding increases
    
    

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

  Across a Period: Ionisation Energy Increases	Down a Group: Ionisation Energy Decreases
  Increase in nuclear charge	Increase in nuclear charge
  Shell number is the same  The distance of outer electrons to the nucleus is the same	Increase in shells Distance of outer electron to nucleus increases The shielding effect increases, therefore, the attraction of outer electrons to the nucleus decreases
  Shielding remains reasonably constant	Increased shielding
  Deceased atomic/ionic radius	Increases atomic/ionic radius
  The outer electron is held more tightly to the nucleus so it gets harder to remove it	The outer electron is held more loosely to the nucleus so it gets easier to remove it

Electron affinity

• When atoms gain electrons they become negative ions or anions
• Electron affinity (EA) can be thought of as the opposite process of ionisation energy and is defined as
  • 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)
• The first electron affinity is always exothermic, e.g. 

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

• 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

Trends in electron affinity

Graph to show the electron affinities across a period

Graph to show the electron affinities from lithium to chlorine

• Electron affinities show periodicity
• The pattern is very similar to ionisation energies, except that it is inverted and the minimum points are displaced one element to the right
• As might be expected, the most exothermic electron affinities are for Group 17 elements which also have the highest electronegativities
• The strongest pull on electrons correlates with the greater amount of energy released when negative ions are formed
• Noble gases do not form negative ions, so they don't appear in this chart
• The electron affinities reach a peak for Group 2 and Group 5 elements

Graph to show the electron affinities down a group

Electron affinities down Group 17 from F to I

• Electron affinities generally decrease down a group
• As the atoms become larger the attraction for an additional electron is less, since the effective nuclear charge is reduced due to increased shielding
• Electron affinity become less exothermic going down the group
• An exception to this is fluorine whose electron affinity is smaller than expected
• This is because fluorine is such a small atom and an additional electron in the 2p subshell experiences considerable repulsion with the other valence electrons

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

Across a period

• Electronegativity increases across a period
• The 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
• The nucleus has an increasingly strong attraction for the bonding pair of electrons of atoms across the period
• This results in smaller atomic radii

Diagram to show the trend in electronegativity across a period

Electronegativity increases going across the periods of the periodic table

Down a group

• There is a decrease in electronegativity going 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, there is a decrease in attraction between the nucleus and outer bonding electrons
• We say the effective nuclear charge has decreased down the group

Diagram to show the trend in electronegativities down a group

Electronegativity decreases going down the groups of the periodic table

Table of trends down a group & across a period