Learning Outcomes Emission and absorption spectra of the hydrogen atom. Balmer series in the emission spectrum as an example. Line spectra as evidence for energy levels. Energy sub-levels. Viewing of emission spectra of elements using a spectroscope or a spectrometer.
Atomic structure Spectra
Spectrscope In a light spectroscope, light is focused into a thin beam of parallel rays by a lens, and then passed through a prism or diffraction grating that separates the light into a frequency spectrum.
Continuous Spectrum Emission Spectra
Continuous spectrum A Spectrum in which all wavelengths are present between certain limits.
Emission Sprectrum
Emission spectrum
Spectrum lines When light from an unknown source is analyzed in a spectroscope, the different patterns of bright lines in the spectrum reveal which elements emitted the light. Such a pattern is called an emission spectrum.
Absorption spectrum
Emission Spectrum Shows that atoms can emit only specific energies (discrete wavelengths, discrete frequencies)
hypothesis: if atoms emit only discrete wavelengths, maybe atoms can have only discrete energies
A turtle sitting on a staircase can take on only certain discrete energies
energy is required to move the turtle (electron) up the steps (energy levels) (absorption) energy is released when the turtle (electron) moves down the steps (energy levels) (emission)
energy staircase diagram for atomic hydrogen
bottom step is called the ground state higher steps are called excited states
Balmer Series Balmer analysed the hydrogen spectrum and found that hydrogen emitted four bands of light within the visible spectrum: Wavelength (nm)Color 656.2red 486.1blue 434.0blue-violet 410.1violet
Flame Test Flame Test The following metals emit certain colours of light when their atoms are excited. Metal Colour Sodium (Na)Yellow Lithium (Li)Pink/Red Potassium (K)Purple Copper (Cu)Green Calcium (Ca)Pink Barium (Ba)Yellow/Orange Strontium (Sr)Red/Orange
Learning Outcomes Energy levels in atoms. Organisation of particles in atoms of elements nos. 1–20 (numbers of electrons in each main energy level). Classification of the first twenty elements in the periodic table on the basis of the number of outer electrons.
Bohr
Bohr’s theory Electrons revolve around nucleus in orbits Electron in orbit has a fixed amount of energy Orbits called energy levels If electron stays in level it neither gains nor loses energy
Bohr Atom absorbs energy Electron jumps to higher level Atom unstable at higher levels. Electron falls back to a lower level Atom loses or emits energy of a particular frequency.
quantisation Electrons can have only certain particular values of energy
EVIDENCE FOR ENERGY LEVELS In Hydrogen electron in lowest (n=1) level; ground state Energy given; electron jumps to higher level excited state Falls back and emits a definite amount of energy Energy appears as a line of a particular colour
colours Energy emitted depends on the jumps Different jumps emit different amounts of energy and hence different colours
Main energy levels (shells) Spectroscopic notation for shells. N shell name 1 = K 2 = L 3 = M 4 = N
Bohr Diagram
Bohr Diagrams To draw Bohr Diagrams: 1.Draw the nucleus as a solid circle. 2.Put the number of protons (atomic number) in the nucleus with the number of neutrons (atomic mass – atomic number) under it. 3.Place the number of electrons (same as protons) in orbits around the nucleus by drawing circles around the nucleus. Remember, 1 st shell – 2 electrons, 2 nd shell – 8 electrons, 3 rd shell – 8 electrons, 4 th shell – 18 electrons.
Valency & Groups
Valencies
Atomic structure 2
Learning Outcomes Energy sub-levels. Heisenberg uncertainty principle. Wave nature of the electron. (Non-mathematical treatment in both cases.) Atomic orbitals. Shapes of s and p orbitals. Building up of electronic structure of the first 36 elements. Electronic configurations of ions of s- and p-block elements only. Arrangement of electrons in individual orbitals of p-block atoms.
Heisenberg We cannot know both the position and speed of an electron Therefore we cannot describe how an electron moves in an atom
Einstein..
De Broglie Matter has wave characteristics
2-slit expt..
Expected Result if light and electrons are particles :
Actual result for light and electrons – demonstrates their wavelike nature :
Orbital A region in space where the probability of finding an electron of a particular is high
Electrons moving
Electron paths
Main levels AND THE NUMBER OF ELECTRONS 1 = 2e 2 = 8e 3 = 18e 4 = 32e
Sub-levels Each main level has sub-levels 1has s sub-level only 2 has s and p sub-levels 3 has s,p and d sub-levels 4 has s,p,d and f sub-levels Energy of sub-levels s p d
1s
2s
2p
3d
Electrons in sub-levels s = 2e p = 6e d = 10e f = 14e
Sub-levels 1 = s(2e) 2 = s(2e) + p(6e) = 8e 3 = s(2e) + p(6e) + d(10e) = 18e
The "p" orbital is dumb belled shaped and each P sub level is made of three "p" orbitals (because the P sub level can hold 6 electrons and every orbital holds 2 electrons)
P-orbitals
Electrons in orbitals S holds 2e 3 p orbitals each holds only 2e 5 d orbitals each holds only 2e
Pauli’s exclusion principle Orbital can only hold 2electrons and these electrons must have opposite spins
Pauli's exclusion principle
Aufbau principle Electrons fill levels in a specific order. 1s 2s 2p 3s 3p 4s 3d 4p
AUFBAU
Hunds rule When filling up the orbitals in a sublevel electrons fill then singly at first.
5 electrons
6 electrons Hund’s rule
Electron Configurations He, 2, helium : 1s 2 Ne, 10, neon: 1s 2 2s 2 2p 6 Ar, 18, argon : 1s 2 2s 2 2p 6 3s 2 3p 6 Kr, 36, krypton : 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6
Exceptions to Electron configuration rules Cr Half-filled orbitals give greater stability 1s 2 2s 2 2p 6 3s 2 3p 6 3d 4 4s 2 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 1 Cu Full 3d sub-level gives greater stability 1s 2 2s 2 2p 6 3s 2 3p 6 3d 9 4s 2 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1
Electron Configurations (ions) F -, 10, Flouride: [1s 2 2s 2 2p 6 ] - Cl -, 18, Chloride : [1s 2 2s 2 2p 6 3s 2 3p 6 ] - Na +, 10, Sodium ion: [1s 2 2s 2 2p 6 ] +