STPM Chemistry Form 6 Notes – Terminology and Concepts: The Electronic Structure of Atoms (Part 1)

by BerryBerryTeacher

in Berry Reference (Notes)

The electronic structure of atoms are as fundamental a concept as it can get for chemistry at STPM level. Unfortunately, it can get a little confusing as it is an extension upon what you have learnt in your SPM days. Students have mentioned that the structure of atoms learnt here is very dissimilar to the ones learnt when they were younger. Berry Berry Easy won’t say that it is totally different but would rather call it an extension of knowledge to be gained.

So get ready to read this first part of the  two-part Berry Berry Easy notes on “The Electronic Structure of Atoms – STPM Chemistry Form 6“. Learn all about the hydrogen spectrum, energy levels, ionisation energy, and its calculation for the ionisation energy for hydrogen. You only need to embrace the concepts and this topic will be berry berry easy. Do memorise the table too as it is a rather popular topic in exams.

STPM Chemistry Form 6 – Terminology and Concepts: The Electronic Structure of Atoms (Part 1)

Spectrum – a display of the components of a beam of radiation.

Hydrogen Spectrum

  1. Hydrogen molecules break up to form hydrogen atoms when hydrogen gas (at low pressure) in a discharge tube that has been passed through by an electrical discharge.
  2. Hydrogen molecules do not emit visible light.
  3. Emission spectrum contains separate sets of lines.
  4. Each line corresponds to a light of a particular frequency / wavelength.
  5. The series of lines is called the Balmer series that consist of four colours lines.
  6. These lines are: 656 nm (red), 486 nm (blue-green), 434 nm (indigo) and 410 nm (violet) – visible to the unaided eyes.
  7. Other sets of lines are: infrared region (Paschen series, Brackett series and Pfund series) and ultraviolet (Lyman series).
  8. Convergence limit – wavelength / frequency at which the converging spectral lines merge together.
    Balmer formula
    : 1/λ = RH (1/22 – 1/n2)
    where λ = wavelength, RH = Rydberg constant, n = 3,4 … ∞
  9. Hydrogen gas also emits light in the ultraviolet and infrared regions of the electromagnetic spectrum.
    Rydberg equation
    : 1/λ = RH (1/n12 – 1/n22)

Berry Important Table of Summary for the Series:

Series n1 n2 Type of electromagnetic radiation
Lyman 1 2, 3, 4 … ∞ Ultraviolet
Balmer 2 3, 4, 5 … ∞ Visible
Paschen 3 4, 5, 6 … ∞ Infrared
Brackett 4 5, 6, 7 … ∞ Infrared
Pfund 5 6, 7, 8 … ∞ Infrared

Electronic Energy Levels

  1. The electrons in an atom can exist at certain energy level.
  2. The electron nearest to the nucleus has the lowest energy.
  3. The further the electron from the nucleus, the higher the energy.
  4. Excited state: sufficient energy (heating or electricity discharge) is needed to promote an electron from a lower energy level to higher one.
  5. The electron will not remain at the high energy level because it is unstable.
  6. Therefore, it will fall back to the level from it started or to the intermediate level.
  7. It will lose an amount of energy (energy = difference between the two energy levels).
  8. Convergence of the spectral lines – difference between successive energy levels becomes smaller with the increasing distance of the energy levels from the nucleus.
  9. Quantum radiationsmall amount of radiation emitted by an electron when it falls from higher to a lower energy level.

Planck’s equation: Δ E = E1 – E2 = hf = h (c / λ)
where λ = wavelength, h = Planck’s constant = 6.63 X 10-34 Js,
c = speed of light = 3.00 X 108 ms-1
Δ E = difference in energy of two energy levels (quantum of radiation)

First ionisation energy – minimum energy required to remove one mole of electrons from one mole of the atoms of an element in the gaseous state.

Example: M –> M+ + e

Ground state – is the energy level nearest to the nucleus and it is the lowest possible energy state.

Calculation for the Ionisation Energy of Hydrogen

(Convergence limit in Balmer serier)

  • Step 1: Find frequency difference between the successive lines in the series.
  • Step 2: Plot a graph of frequency difference (y-axis) against the lower frequency (x-axis).
  • Step 3: Extrapolate the graph to obtain the frequency (x-axis) when frequency difference = 0 (convergence limit)
  • Step 4: Calculate the ionisation energy by using ΔE = hf

In the next part, you’ll learn the very important subsection on electronic configurations.

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