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 2025 SL- IB Style Practice Questions with Answer-Topic Wise-Paper 1
- IB DP Chemistry 2025 SL- IB Style Practice Questions with Answer-Topic Wise-Paper 2
- IB DP Chemistry 2025 HL- IB Style Practice Questions with Answer-Topic Wise-Paper 1
- IB DP Chemistry 2025 HL- IB Style Practice Questions with Answer-Topic Wise-Paper 2
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
- When the electron moves up an energy level it absorbs energy
- 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)