Chemistry is in the electrons Electronic structure – how the electrons are arranged inside the atom Two parameters: –Energy –Position

Learning objectives Describe the properties of waves Describe the photon and the relationship between light frequency and energy 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

Why don’t the electrons collapse into the nucleus? Like charges attract

The planetary idea Unlike the moon, an orbiting charged particle will spiral inwards to the nucleus Moving charges lose energy by radiation

Conventional explanations do not work Yet we know that atoms exist and that electrons are stable outside the nucleus. Another explanation must be found The iconic symbol of the planetary atom is not correct

Let there be light Electromagnetic waves are characterized by: –Wavelength – distance between two peaks –Frequency - number of wave peaks that pass a fixed point per unit time –Amplitude - height of the peak measured from the center line Different “colours” of electromagnetic radiation are waves with different wavelengths and frequencies.

Wavelength and frequency are related All electromagnetic radiation has the same velocity: the speed of light – 3 x 10 8 m/s Velocity = wavelength x frequency Wavelength proportional to 1/frequency - as wavelength ↑ frequency ↓

Continuous range of wavelengths and frequencies: –Radio waves at low-frequency end –Gamma rays at high-frequency end –Visible region is a small slice near the middle Note X-rays have wavelength approximately same as diameter of an atom (10 –10 m) Radiation becomes more dangerous as frequency increases Harmful radiation The electromagnetic spectrum

Atoms emit and absorb radiation at specific wavelengths Absorption is light removed by the atom from incident light Emission is light given out by an energetically excited atom The absorption and emission lines are at the same wavelengths The lines from the H atom form a neat series. Do these spectra have anything to do with the electronic structure?

Each element has a unique spectrum Electronic structures of each element are different Spectra can be used to identify elements – even in very remote locations

Light is a wave and a particle – the photoelectric effect Light incident on a metal surface causes electrons to be emitted. Below threshold frequency nothing happens Above threshold, current increases with intensity No current flows Current flows

Quantization: light as particle and wave In photoelectric effect, light behaves like a particle. Energy is “quantized” into packets Each packet is a photon Photon energy depends on frequency: E = h E = h ν As frequency increases photon energy increases

Photon energy increases from right to left “Dangerous” radiation has high photon energy – UV light, X-rays Harmless radiation has low photon energy – IR, radiowaves Blue light has higher energy than red light

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

Successes and shortcomings of Bohr Could not explain why these levels 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

Standing waves and strings Strings of fixed length can only support certain wavelengths. These are standing waves.

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

Orbits become orbitals

Orbitals are described by quantum numbers Each orbital has unique set 1s, 2p, 3d etc. Number (Principal quantum number) describes energy Letter describes shape –s zero dimensions –p one dimension –d two dimensions –f three dimensions

Levels and sublevels Levels are distinguished by principal quantum number n (n = 1, 2, 3...) Sublevels are all members of same level (same n) but have different letters (s, p, d, f) Number of different sublevels in level n = n