Topic 7: Periodic Trends

T7VA: Atomic Size

3 periodic trends to keep track of

  1. atomic/ionic size

  2. ionisation energy

  3. electron affinity

Trends show us how physical properties change as we move left to right, bottom to top in the periodic table

Atomic size: the radius of an atom

  • measure the bond length and divide by 2

Atomic size trend: atomic size increases from R → L and U → D

Atomic size can be understood through Coulombs Law:

Coulombs Law describes the force of attraction and repulsion between two charges (electrostatic force).

q1 and q2 refer to the two charges

  • If the two charges are equal and positive, they will repell

  • if the two charges are equal and negative, they will attract

  • if one charge is 3x as positive and the other is 1x as negative, the positive charge will attract the negative one with a much stronger pull

As we move to the right on the periodic table, protons attract the electrons more. As the number of protons in an atom increase, the radius decreases because the attractive pull of the many protons on the electrons is stronger the more protons there are. However, the radius of the atom also is affected by the electrons’ repelling as they are moved close and closer by the attraction of the protons. Therefore, the radius of an atom is the result of a balance between the attractive force between the protons and electrons, and the repulsive force between the electrons.

As we go down the periodic table, we increase the number of atomic orbital energy levels that surround the nucleus, and so the overall size of the atom increases.

T7VB: Ionic Size

Ions increase in size the more negative their charges are

  • Cations are always SMALLER than the atoms they come from

  • Anions are always LARGER than the atoms they come from

as

REMEMBER: an ion can be isoelectronic with an atom if they have the same amount of electrons

  • For example, K+1 and Cl-1 are isoelectronic with Ar

T7VC: Ionisation Energy Part 1

Ionisation energy is the energy it takes to remove an electron from an atom

  • ionisation energy increases as you keep removing electrons (second ionisation energy is greater than first ionisation energy, etc)

    • first ionisation: removal of the first electron, notated as I1

    • second ionisation: removal of the second electron, notated as I2; etc…

as you remove electrons, ionisation energy increases because the atomic radius decreases due to there being less and less electron repulsion

  • fewer electron-electron repulsions result in overall lower energy, so as ionisation continue, the energy of the inner shells is lower and lower

Where there is a big jump in ionisation energy, you can tell the ionisation process has moved from stripping valence electrons to stripping core electrons

  • Inner shell electrons: core electrons

  • Outer shell electrons: Valence electrons

Trend: ionisation energy increases moving L → R, D → U

  • IE increases in a period moving Left to Right

NOTE: there are some exceptions

T7VD: Ionisation Energy Part 2

There are some exceptions to the general trend of IE increasing L → R, D → U

Exception 1:

Why is the ionisation energy for Mg higher than that for Al ?

  • Al is more to the right side of the periodic table than Mg, right ?

this is because of their electron configurations:

  • Mg: [Ne] 3s2

  • Al: [Ne] 3s23p1

The 3p orbital is higher in energy than 3s, so it is easier to remove electrons from it. The electrons in the 3p orbitals are shielded from the nucleus by the 3s orbital…

  • shielded electrons are less strongly attracted to the nucleus and are thus more easily removed

Therefore, the 3p1 electron is easier to ionise than the 3s2

Exception 2:

Why is the ionisation energy for P higher than that for S ?

this is ALSO because of their electron configurations:

p: three unpaired electrons that all have the same spin

s: four electrons, two of which are paired, and three of which are unpaired with the same spin

NOTE: paired electrons are HIGHER ENERGY than unpaired electrons within the same suborbital

Therefore…

  • the Sulphur p electrons are at a higher energy than the Phosphorus p electrons

  • the Sulfur p orbital is at a higher energy than the Phosphorus p orbital

  • electron-electron repulsions in Sulphur are higher than those in Phosphorus because S’s electrons are closer together

SO, the 4th p electron in Sulfur is easier to remove and thus has a lower first ionisation energy than Phosphorus

Overview:

There are some exceptions to the general trend of IE increasing L → R, D → U

These exceptions are:

  1. the electron - electron repulsion in the 4th p electron as you move from group 15 to 16 makes the 4th p electrons easier to remove than the 3rd p electrons

  2. the shielding of the p electrons in group 13 by their s orbital electrons makes it so that the 1st p electrons are easier to remove than the 2nd s electron

T7VE: Electron Affinity

Electron affinity (EA): the readiness to accept electrons

  • the heat released when a gaseous atom gains an electron to create a negative ion

  • X(g) + e- → X-(g)

EA = -ΔH

  • -ΔH means that energy is released: exothermic

  • the more negative the ΔH value, the more heat is released, and the larger the electron affinity

  • if a lot of energy lost in the formation of the ion means that the ion is more stable than the atom

Group 17 atoms (halogens) have the largest electron affinities of their respective periods. Why ?

  • the halogens all have an eleconfig ending in 2p5. This means that it is more favourable for them to gain an electron and fill up that last shell than lose electrons.

  • However, as you move down the period, there are more and more shells, which means that the outermost 2p5 electrons are further and further away from the nucleus and so have less attraction to the protons of the nucleus. Therefore, the binding energy is weaker and so less heat is released in the process of gaining an electron

EXAMPLE: Cl

In the process of Cl gaining an electron, 349 kJ of energy are released. Therefore, its ΔH is -349 kJ/mol. That means the EA is |-349| = +349 kJ/mol.

This is a favourable process, and the Cl-1 ion is more stable compared to the Cl atom.

Trend: EA affinity increases from L → R, D → U

  • NOTICE:

    • halogens have the largest electron affinities

    • the metals have very small EA compared to the halogens because forming anions is not favourable

EXEPTIONS: atoms in the p block of period 2 have EA that are somewhat lower than the atoms in period 3

  • this this is because the electrons of the p block of period 2 are really small, and so the electron-electron repulsion is greater than in period 3

  • noble gases (and a few others) have no EA values listed because the formation of anions is not favourable either

Lower electron affinity = larger atomic size = lower ionisation energy

Electronegativity

Electronegativity increases going L → R, D → U