1. Light Light is a kind of electromagnetic radiation, which is a from of energy that exhibits wavelike behavior as it travels through space. Other forms of electromagnetic radiation include X-rays, UV rays, infrared rays, microwaves and radio waves. All forms move at a constant speed of 3.0 x 10 8 m/s

2. Wavelength and frequency Wavelength is the distance between corresponding points on adjacent waves. Frequency is the number of waves that pass a given point in a specific time (usually 1 second).

3. The photoelectric effect The photoelectric effect refers to the emission of electrons from a metal when light shines on the metal. If the light’s frequency was below a certain minimum no electrons were emitted, no matter how long it was shone. Planck suggested that the light was emitted as quanta of energy rather than waves. E = hv

4. Photons Einstein – electromagnetic radiation has a dual wave-particle nature. Each particle of light carries a quantum of energy, called a photon.

5. The lowest energy state of an atom is its ground state. A state in which an atom has a higher potential energy than its ground state is called an excited state. When an excited atom returns to its ground state, it gives off energy it gained in the form of electromagnetic radiation.

6. Emission Spectra Emission spectra – produced by passing electric current through a vacuum tube containing gas at low pressures, and then shone through a prism. Not continuous. Hydrogen emits only at specific frequency – energy differences between the atoms’ energy states are fixed.

7. Bohr’s Model Neils Bohr proposed a model that linked the atom’s electron with photon emission. He said that electrons could only orbit in fixed energy states. His model explained only the hydrogen atom.

8. Quantum model of the atom De Broglie – electrons could also be considered as waves. Electron waves could exist only at specific frequencies. Heisenberg uncertainty principle – it is impossible to determine at the same time, both the position and velocity of an electron or any other particle Quantum theory describes mathematically the wave properties of electrons and other very small particles.

9. Orbitals Wave functions give only the probability of finding an electron at a given place around the nucleus Do not travel in neat orbits, but exist in certain regions called orbitals. An orbital is the three dimensional region around the nucleus that indicates the probable location of an electron.

10. Quantum numbers The principal quantum number (n) indicates the main energy level occupied by an electron

11. The angular momentum quantum number (l) indicates the shape of the orbital. s is spherical p is dumb-bell shape d is complex

12. Magnetic quantum number indicates the orientation of an orbital around the nucleus. Spin quantum number – only two possible values

13. Electron Configuration Aufbau’s principle

14. Bohr Model – energy levels An electron can be in any of these energy levels, but cannot be found in between.

15. In the ground state – atom is unexcited – electrons are in the lowest possible energy level. If an atom is excited the electron may go up to the next level. When they fall back they emit light. This is an element ’ s emission spectrum

16. Emission spectra

17. Later complications – De Broglie – electrons act both as particles and waves Heisenberg – you can never know exactly the position and velocity of an electron (or any other particle) Schrodinger – wave equations – (quantum theory) An orbital is the region around the nucleus where an electron can be found.

18. Why do we need to know where electrons are found in the atom? Because the chemical behavior (reactivity) of an element depends on the energy levels of its electrons

19. The Principal quantum number, n, is the main energy level occupied by an electron

20. Sublevels – the shape of the orbital

21. d-orbitals

22. Spin quantum number A single orbital can hold a maximum of two electrons with opposite spins

23. Filling electron orbitals

25. Filling energy levels (Aufbau ’ s rule)