STPM Chemistry Form 6 Notes – Terminology and Concepts: The Periodic Table (Part 3)

by BerryBerryTeacher

in Berry Reference (Notes)

This post is Part 3 of the “STPM Chemistry Form 6 notes on The Periodic Table” from Berry Berry Easy which focuses on arguably the trickiest part in the periodic table. Here, you’ll learn all about the three factors (covered in STPM Chemistry syllabus) which affects the atomic radii, namely the screening effect, nuclear charge and effective nuclear charge. Using that information, STPM students should be able to explain the change in atomic radius across and down a group. (This is a popular topic in exams of all levels) Similarly, you’ll also learn about the ionic radius. So make sure you can distinguish between the two radii.

It is crucial that you understand this post fully because every question about the periodic table will invariably be linked to the content from this post..

STPM Chemistry Form 6 Notes – Terminology and Concepts: The Periodic Table (Part 3)

Periodicity of Atomic Radius

Atomic radii for elements in Periods 2 and 3

Elements Atomic radius (pm)
Li 152
Be 112
B 80
C 77
N 74
O 74
F 72
Na 156
Mg 136
Al 125
P 110
S 104
Cl 99

Atomic radii can be classified into three categories:

  • Covalent radius:
  • Metallic radius
  • Van der Waals radius

Effecting factors of the atomic radius:

  • Screening effect of the inner shell electrons: negatively-charged shells repel one another and are being pushed further away from the nucleus; screening effect increase; and size of the atoms increase.
  • Nuclear charge (number of protons in the nucleus) that pulls all the electrons closer to the nucleus: The higher the nuclear charge; the stronger the attraction between nucleus and the electron cloud; and the size of the atom decrease.
  • Effective nuclear charge = No. of protons – No. of inner electrons

A) Atomic radius across a period

Example: Period 2 (Li, Be, B, C, N, O, F, Ne) and Period 3 (Na, Mg, Al, Si, P, S, Cl, Ar)

Across the period:

  • Number of protons increase by one.
  • Number of electrons increase by one.
  • Screening effect does not affect much (same quantum shell).
  • Nuclear charge increase (stronger attraction between nucleus and electron cloud).
  • Size of the atoms decrease.

B) Atomic radius down a group

Example: Group 2 (Be, Mg, Ca, Sr, Ba)

Down the group:

  • Screening effect increase.
  • Nuclear charge increase.
  • Effective nuclear charge decrease.
  • Size of the atoms increase (the increase in the screening effect is larger than the increase in the nuclear charge).

C) Ionic radius (radius of a cation or or an anion) across Period 3

Ion Ionic radius No. of electrons No. of protons
Na+ 0.095 10 11
Mg2+ 0.065 10 12
Al3+ 0.050 10 13
P3- 0.212 18 15
S2- 0.184 18 16
Cl- 0.181 18 17

Isoelectronic – species have the same number of electrons and the same electronic configuration.

When given number of electrons (Na+, Mg2+, Al3+) or (P3-, S2-, Cl-)

  • higher the nuclear charge,
  • higher the force of attraction
  • smaller the atomic size or ionic size.

When given nuclear charge,

  • larger the number of electrons in an atom or an ion,
  • greater the repulsion between electrons
  • larger the atomic or ionic size.

Conclusion:

  • Cationic size decreases (increasing proton number).
  • Anionic size decreases (increasing proton number).

D) Ionic radius down a group

Example: Group 2 (Be2+, Mg2+, Ca2+, Sr2+, Ba2+) & Group 17 (F-, Cl-, Br-, I-)

Going down the Group 2 and Group 17:

  • Each successive ion has one additional shell filled with electrons.
  • Screening effect increase
  • Ionic size increase.

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