MM 2006 AH Spectra Continuous spectrum Emission spectrum low pressure gas emitted light Prism or diffraction grating

MM 2006 AH Emission spectrum of neon, nm

MM 2006 AH Emission spectrum of hydrogen The Balmer series , nm E = Lh c = energy emitted when excited electrons fall back to a lower energy level frequency of line due to difference in energy between 2 electron energy levels E = L h c

MM 2006 AH Energy levels E3E3 E2E2 E1E1  E b = E 3 - E 2 = h b  E a = E 3 - E 1 = h a frequency of lines in emission spectrum fixed energy between levels fixed energy of electrons fixed - quantised - photon emitted when electron moves to lower energy level + E1E1 E2E2 E3E3

MM 2006 AH Full emission spectrum of hydrogen SeriesEnergy level excited electron falls toPart of electromagnetic spectrum Lymann = 1Ultra-violet Balmern = 2Visible Paschenn = 3Infra-red Brackettn = 4Infra-red Pfundn = 5Infra-red P7 & 8 LTS

MM 2006 AH The Lyman series , nm convergence limit p9 LTS Q

MM 2006 AH Bohr’s theory the hydrogen atom’s electron exists only in certain definite energy levels the electron changes energy levels when a photon is absorbed or emitted the energy of the photon equals the difference between the two energy levels  E = h. Quanta - fixed quantities of energy possessed by electrons

MM 2006 AH Quantum mechanics electrons - wave-particle duality 4 quantum numbers define energy of an electron electrons arranged in shells shells described by principal quantum number, n n = 1, closest to nucleus; second shell n = 2 the higher n, the higher the potential energy associated the shell and the further from the nucleus the electron likely to be found Spectra show doublets and triplets suggesting subshells s, p, d, f

MM 2006 AH Shells and subshells ShellSubshells 11s 22s, 2p 33s, 3p, 3d 44s, 4p, 4d, 4f subshells have different energies s < p < d < f

MM 2006 AH Second quantum number each subshell contains one or more energy levels or orbitals orbitals defined by angular momentum quantum number, l l related to shape of orbital l value 0, 1, 2 …… (n -1) Heisenberg’s uncertainty principle governs behaviour of electrons impossible to define a point in space where electron certain to be found regions in space where probability of finding an electron high - atomic orbitals n defines orbital size l defines shape of orbital

MM 2006 AH s orbitals 1s2s3s

MM 2006 AH Third quantum number m 1 magnetic quantum number gives number of orbitals and spatial orientation m 1 any integer between -l and +l defines p, d, f orbitals

MM 2006 AH p orbitals l = 1 3 possible values of m 1, -1, 0, +1 3 orbitals of equal energy - degenerate arranged along 3 axes x, y, z e.g. 2p x, 2p y, 2p z

MM 2006 AH d orbitals l = 2 m 1 = -2, -1, 0, +1, +2 hence 5 d orbitals f orbitals l = 3 m 1 = -3, -2, -1, 0, +1, +2, |+3 hence 7 f orbitals

MM 2006 AH Fourth quantum number electrons spin spin quantum number, m s + 1/2 and - 1/2