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Electron configurations – IB DP Chemistry- Study Notes | IITian Academy

Electron configurations - IB DP Chemistry- Study Notes - New Syllabus 2025

Electron configurations – IB DP Chemistry- Study Notes

IITian Academy excellent Introduction to the Particulate Nature of Matter – Study Notes and effective strategies will help you prepare for your IB DP Chemistry  exam.

IB DP Chemistry Study Notes – All Topics

S1.3.1 – Emission Spectra

Types of Spectrums:

  • Emission Spectra
  • Warm gas
  • Absorption Spectra
  • Cold gas

Continuous Spectrum: A type of absorption spectrum

  • This spectrum contains all types of wavelengths
  • A continuous series of colors with each color merging into the next
  • There are no gaps in between the colors
  • Wavelength of Visible Light: 400nm – 700nm
  • An example of a continuous spectrum is a rainbow. 

Line Spectrum: A type of emission spectrum

  • Sharp lines produced with specific frequencies
  • Produced by excited atoms and ions as they fall back to a lower energy level.
  • Different elements have different line spectra so they can be used to identify unknown elements.

The EM Spectrum:

  • Electromagnetic radiation comes in different forms
  • All forms travel at the same speed of light but with different wavelengths
  • Higher energy forms have shorter wavelengths and higher frequencies.

Wavelength, Frequency, and Energy: A photon is a quantum of energy, and it is proportional to frequency (f) and radiation.

E: energy possessed by a photon in Joules (J)

h: Planck’s Constant, 6.63 x 10-34 Js

f: frequency of radiation, this is in Hertz (Hz) or s-

c= speed of light

c = 3 x 108

(Frequency is also (v))

THE EQUATION:

E = hf

   OR

E = h(c/λ)

f  or v = (c/λ)

S1.3.2-3 – Line Emission Spectra of Hydrogen

The Hydrogen Spectrum: The hydrogen atom gives out energy when an electron fall from a higher to a lower energy level.

The hydrogen Spectrum has 4 major characteristic lines Violet, blue, green, and red

  • When an electron falls to the third energy level(n=3) or higher, infrared (IR) radiation is produced
  • Hydrogen produces visible light when an electron falls to the second energy level(n=2)
  • Transition to the first energy level (n=1)is a high energy change, this is the ultraviolet (UV) region of the spectrum  
  • The HIGHER the energy change, the SHORTER the wavelength
  • The HIGHER the energy change, the HIGHER the frequency

Electrons and Energy Level:

  • The maximum number of electrons in any energy level (n) is 2n2
  • Each energy level is divided into 4 sub-levels:
  • ‘s’ orbital: 2 electrons
  • ‘p’ orbital: 6 electrons
  • ‘d’ orbital: 10 electrons
  • ‘f’ orbital: 14 electrons
  • As the shells increase, the closer the energy is together
  • As shown on the side, the energy levels get closer together as the shells increase

Pauli Exclusion Principle:

  • No more than 2 electrons occupy an orbital

Hund’s Rule:

  • Orbitals are filled individually first

Aufbau Principle:

  • Electrons fill low-energy orbitals first

-the visible region is from n=2 until n=7 (Balmer series)

-n=1 has the highest energy (uv) any transition that ends at n=1 has high energy (lyman series)

-if the electron lands on n=3 or above then it emits infrared radiation(Paschen and Bracket series)

Bohr’s theory

  1. When the electron moves up an energy level it absorbs energy
  2. When it moves down it emits energy

Limitations of Bohr

-It is not possible to find the position and momentum of electrons

n

2n2

Maximum Electrons

1

2X12

2

2

2X22

8

3

2X32

18

4

2X42

32

5

2X52

?

Exceptions:

Cr24 – 1s2/2s22p6/3s23p6/3d54s1

Cu29 – 1s2/2s22p6/3s23p6/3d104s1

S1.3.5 – Electron Configuration

Clarification:

  • To clarify, electron arrangement gives the number of electrons in each main energy level.
  • Electron configuration means the number of electrons in each sub-level.

Energy Level

Number of Sub-level

(orbitals)

sublevels

Number of electrons in each sublevel

Total number of electrons in each

n=1

1

1s

2

2

n=2

2

2s,2p

2,6

8

n=3

3

3s,3p,3d

2,6,10

18

n=4

4

4s,4p,4d,4f

2,6,10,14

32

S3.1- [HL] Electron Configuration

First Ionisation Energy:

  • First ionization energy is the energy (per mole) needed to remove one electron from an atom in a gaseous state
  • Measured in kg/mol
  • X(g)→ X+(g) + e-

Successive Ionisation Energy:

  • Increase because electrons are more difficult to remove as they get closer to the atom because the attraction becomes stronger.

General information about Electron Configuration and Ionisation Energy:

  • Ionization energy INCREASE from LEFT TO RIGHTon the periodic table
  • Going left to right, the nuclear energy increases, making it more DIFFICULT to remove an electron
  • Ionization energy DECREASES from UP TO DOWNon the periodic table
  • Why?
  • Outermost electrons are further from the nucleus

ATOMIC RADIUS: This is NOT a part of S3.1 but for info.

  • Atomic radius DECREASES from LEFT TO RIGHTon the periodic table
  • Atomic radius INCREASES from UP TO DOWNon the periodic table.

Ionization Energy Exceptions:

  • Boron (B) and Beryllium (Be)
  • 5B = 1s2/2s22p1
  • 4Be = 1s2/2s2
  • Be > B in terms of ionization energy
  • The (2p1) is easier to remove in Boron. The (2p) orbital electron is farther from the nucleus which makes the electrostatic attraction weaker as compared to (2s) electrons
  • Magnesium (Mg) and Aluminium (Ai)
  • 12Mg = 1s2/2s22p6/3s2
  • 13Ai = 1s2/2s22p6/3s23p1
  • Mg > Ai
  • The (3p1) is easier to remove in Aluminium because these electrons will be better used to shield the (3s) electrons instead of shielding themselves.
  • Oxygen (O) and Nitrogen (N)
  • 😯 = 1s2/2s22p4
  • 7N = 1s2/2s22p3
  • O < N
  • The (2s) orbital is easier to remove in Oxygen,
  • Nitrogen already has the maximum unpaired electrons and is at the lowest energy – least amount of electrostatic repulsion.
  • Phosphorus (P) and Sulfur (S)
  • 15P = 1s2/2s22p6/3s23p3
  • 16S = 1s2/2s22p6/3s23p4
  • P > S
  • Less energy required to remove an electron from (3p) orbital or Sulfur as compared to Phosphorus (the 3p orbitals of it are half-filled)
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