Can be described in terms of waves and characterised in terms of wavelength and/or frequency
Relationship between wavelength and frequency
c = f λ
Electromagnetic spectrum
Different types of radiation arranged in order of wavelength
Wavelengths of visible light are normally expressed in nanometres (nm)
Electromagnetic radiation
Can be described as a wave (has a wavelength and frequency), and as a particle, and is said to have a dual nature
Photon
Particles of electromagnetic radiation
Photon energy
Proportional to the frequency of radiation
Photon energy
Higher frequency radiation can transfer greater amounts of energy than lower frequency radiation
Energy of a single photon
E = hf = hc/λ
Energy of one mole of photons
E = Lhf = Lhc/λ
Energy is often in units of kJ mol-1
Absorption of energy by atoms
Electrons within the atoms may be promoted to higher energy levels
Emission of light energy by atoms
Excited electrons move from a higher energy level to a lower energy level
The light energy emitted by an atom produces a spectrum that is made up of a series of lines at discrete (quantised) energy levels
Each element in a sample produces characteristic absorption and emission spectra
Absorption spectroscopy
Electromagnetic radiation is directed at an atomised sample, and radiation is absorbed as electrons are promoted to higher energy levels
Absorption spectrum
Produced by measuring how the intensity of absorbed light varies with wavelength
Emission spectroscopy
High temperatures are used to excite the electrons within atoms, and as the electrons drop to lower energy levels, photons are emitted
Emission spectrum
Produced by measuring the intensity of light emitted at different wavelengths
Atomic spectroscopy
The concentration of an element within a sample is related to the intensity of light emitted or absorbed
The discrete lines observed in atomic spectra can be explained if electrons, like photons, also display the properties of both particles and waves
Atomic orbitals
Electrons behave as standing (stationary) waves in an atom, and there are different sizes and shapes of standing wave possible around the nucleus, known as orbitals
Orbitals can hold a maximum of two electrons
Types of orbitals
s, p, d and f (knowledge of the shape of f orbitals is not required)
Quantum numbers
Principal quantum number n, angular momentum quantum number l, magnetic quantum number ml, spin magnetic quantum number ms
Principles governing electron arrangement in atoms
Aufbau principle, Hund's rule, Pauli exclusion principle
In an isolated atom the orbitals within each subshell are degenerate
Orbital box notation
Used to represent the relative energies of orbitals in a multi-electron atom
Electronic configurations using spectroscopic notation and orbital box notation can be written for elements of atomic numbers 1 to 36
Periodic table blocks
s, p, d and f blocks correspond to the outer electronic configurations of the elements within these blocks
Variation in ionisation energies
Can be explained in terms of the relative stability of different subshell electronic configurations
There is a special stability associated with half-filled and full subshells
VSEPR theory
Used to predict the shapes of molecules and polyatomic ions
Determining the number of electron pairs surrounding a central atom
Take the total number of valence electrons, add one for each atom attached, add an electron for every negative charge, remove an electron for every positive charge, divide by two
Arrangements of electron pairs around a central atom