Topic 2.2: Electron configuration Flashcards
Electromagnetic spectrum
Spectrum of wavelengths that comprise the various types of electromagnetic radiation
Relation of energy, wavelength, and frequency
a) Energy is inversely proportional to wavelength
b) Energy is proportional to frequency
Electromagnetic radiation
A form of energy that propagates through space at the speed of light as photons
Continuous spectrum
Radiation that spreads all frequencies / wavelengths of light present
Line spectrum
Radiation that emits only certain frequencies/wavelengths of light present
Line spectrum and element
Each element has its own characteristic line spectrum which can be used to identify the element
Quantization
Electromagnetic radiation comes in discrete parcels.
Energy equation
E = hv = hc / A
Electron transition in energy levels
a) Electron can move to a higher energy level by absorption of a photon
b) Electron can move from an excited state to a lower energy level by emitting a photon
Hydrogen line spectrum
Discrete lines which converge at higher energies and form a continuum
Ionization in line spectrum
Beyond the convergence limit the electron can have any energy and is not longer in the atom.
Series of lines in the hydrogen line emission spectrum
a) Balmer series
b) Lyman series
c) Paschen series
Balmer series
a) Visible region
b) Electronic transitions from upper energy levels back down to the n = 2 energy level.
Lyman series
a) UV region
b) Electronic transitions from upper energy levels back down to the n = 1 energy level.
Paschen series
a) IR region
b) Electronic transitions from upper energy levels back down to the n = 3 energy level.
Emission spectrum
Spectrum of frequencies of ER emitted due to an atom making a transition from a high energy to a lower energy state.
How is an emission spectrum formed?
a) Passing an electric discharge through a gas causes an electron to be promoted to a higher energy level (shell).
b) The electron is unstable in this higher level and will fall to a lower energy level, giving out extra energy in the form of a photon light
Electron shell (n = 1,2,3)
a) Main energy level
b) Can hold a maximum number of 2n^2 e-
Subshells (s < p < d < f)
Each subshell has a particular number of orbitals
Atomic orbital
a) Region in space where there is a high probability of finding an electrom
b) Can hold a maximum of two electrons of opposite sign
Electron arrangement
Shows the number of electrons in each shell or orbit
Aurfbau principle
Electrons fill sub-levels from the lowest energy level upwards
Reasons for removing electron from 4s before 3d levels of 3d-block elements
3d orbitals are more compact than the 4s orbitals and hence electrons entering the 3d orbitals will experience a much greater mutual repulsion.
Pauli exclusion principle
Any orbital can hold a maximum of two electrons, having opposite spin
Hund’s rule of maximum multiplicity
When filling degenerate orbitals, electrons fill all the orbitals singly before occupying them in pairs to minimize repulsion
Reason for exceptions in Cr and Cu
Having the maximum number of electron spins the same within a set of degenerate orbitals gives a lower energy (more stable) situation
