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

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

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

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

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

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

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

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

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

Electron microscopes can peer within – waves interacting with matter

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

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

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

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

Orbital energy levels in H and other elements

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

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

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

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

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

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

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

Electrons don’t like to pair C(z = 6)

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

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

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

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

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

The periodic law

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