The Bohr model for the electrons

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1 The Bohr model for the electrons Electronic structure how the electrons are arranged inside the atom Applying the quantum principle of energy Two parameters: Energy Position

2 Learning objectives Describe the basic principles of the Bohr model Distinguish between the classical view and the quantum view of matter Define atomic orbitals Distinguish between the Bohr orbit and atomic orbital Apply quantum numbers and atomic orbitals to building atoms and the periodic table Describe periodic trends in terms of electronic structure

3 Bohr s theory of the atom: applying photons to electronic structure Electrons occupy specific levels (orbits) and no others Orbits have energy and size Larger orbits are at higher energy Electron excited to higher level by absorbing photon Electron relaxes to lower level by emitting photon Photon energy exactly equals gap between levels

4 Size of energy gap determines Small energy gap, low frequency, long wavelength (red shift) High energy gap, high frequency, short wavelength (blue shift) photon energy

5 The full spectrum of lines for H Each set of lines in the H spectrum comes from transitions from all the higher levels to a particular level. The lines in the visible are transitions to the second level

6 The Bohr orbits Bohr orbits have quantum numbers n n = 1 (capacity 2) n = 2 (capacity 8) n = 3 (capacity 8)

7 Bohr orbits and the periodic table Elements in the same group have the same number of electrons in outer Bohr orbit

8 Successes and shortcomings of Bohr Couldn t explain why orbits were allowed Only successful agreement with experiment was with the H atom Introduced connection between spectra and electron structure Concept of allowed orbits is developed further with new knowledge Nonetheless, an important contribution, worthy of the Nobel prize

9 Electrons are waves too! Life at the electron level is very different Key to unlocking the low door to the secret garden of the atom lay in accepting the wave properties of electrons De Broglie wave-particle duality All particles have a wavelength wavelike nature. Significant only for very small particles like electrons or photons As mass increases, wavelength decreases Electrons have wavelengths about the size of an atom Electrons are used for studying matter electron microscopy

10 Electron microscopes can peer within waves interacting with matter

11 Heisenberg Uncertainty Principle: the illusive electron We can predict the motion of a ball; But not an electron: problems locating small objects

12 The Quantum Mechanics: waves of uncertainty System developed that incorporated these concepts and produced an orbital picture of the electrons No longer think of electrons as particles with precise location, but as waves which have probability of being in some region of the atom the orbital Impossible with the classical mechanics of Newton

13 Orbits become orbitals The Orbitron: a gallery of atomic orbitals and molecular orbitals

14 Orbitals are described by quantum numbers Each orbital has unique set 1s, 2p, 3d etc. Number describes energy Letter describes shape S zero dimensions P one dimension D two dimensions F three dimensions

15

16 Getting from the orbitals to the elements All elements have the same set Atomic number dictates how many are filled how many electrons are added Filling orbitals follows a fixed pattern: lowest energy ones first

17 Orbital energy levels in H and other elements

18 How many per orbital? Electrons share orbitals (only two allowed) A consequence of spin

19 How many electrons can be added to the orbitals 1s, 2s, 3s etc. 2 electrons 2p, 3p, 4p etc. 6 electrons 3d, 4d etc. 10 electrons 4f, 5f etc. 14 electrons

20 Add electrons to the orbitals lowest first 4s 4p 3d 3s 3p 2s 1s 2p H(z = 1)

21 Fill lowest orbital 4s 4p 3d 3s 3p 2s 1s 2p He(z = 2)

22 Begin next orbital 4s 4p 3d 3s 3p 2s 1s 2p Li(z = 3)

23 Fill 2s 4s 4p 3d 3s 3p 2s 1s 2p Be(z = 4)

24 Begin filling 2p 4s 4p 3d 3s 3p 2s 1s 2p B(z = 5)

25 Electrons don t like to pair 4s 4p 3d 3s 3p 2s 1s 2p C(z = 6)

26 4s 4p 3d 3s 3p 2s 1s 2p O(z = 8)

27 4s 4p 3d 3s 3p 2s 1s 2p F(z = 9)

28 Filled 2p neon unreactive 4s 4p 3d 3s 3p 2s 1s 2p Ne(z = 10)

29 Shape of the periodic table explained by orbital picture 2 groups 10 groups 6 groups 14 groups

30 Connecting the table with orbitals: elements per row matches capacity of orbitals

31 Simplifying with shells: echoes of Bohr orbits The orbitals with the same Principal Quantum number (1,2,3 etc) are grouped into shells Filled shells have special significance Filled shell 1 Filled shell 1 Unfilled shell 2 Filled shell 2

32 The periodic law and atomic size

33 Ionization energy and the periodic law Ionization energy is energy required to remove electron from the neutral atom

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