1. Review Test Monday April 13 th MC (10 marks) Short answer (37 marks)

2. Review History – Dalton All matter is composed of indivisible particles called ATOMS Atoms of a given element are identical in mass and properties Compounds are formed by a combination of 2 or more atoms Atoms cannot be created or destroyed

3. History J.J Thompson – Using a cathode ray, Thompson discovered the subatomic particle called the ELECTRON were present in atoms Small Negatively charged – PLUM PUDDING MODEL

4. History Rutherford – Conducted gold foil experiment – Sent a positive stream of particles through a piece of gold foil – Predicted that all particles should pass through however, every now and then a particle would deflect back – Concluded: Atoms are mostly empty space Center of the atom is a dense positive area called the NUCLEUS

5. History Bohr Electrons orbit around the nucleus in shells (solar system) 7 different shells or orbits Electrons in different orbits have different energy levels electrons would only occupy the lowest possible energy level Electrons can move up a level if the lower levels were full

6. History Schrödinger – WAVE-MECHANICAL MODEL Electrons move in areas called ORBITALS which is an area of high probability of finding an electron Energy can behave as both WAVES and PARTICLES

7. Light and Energy WAVES – Amplitude- height of the waves from the origin to the crest – Wavelength (λ lambda) – distance between crests – Frequency – (f) is the number of wave cycles to pass in a given period of time

8. Light and Energy Higher Frequency Higher Energy

9. Light and Energy Bohr – Electrons orbiting in shells around the nucleus – Energy levels differ from one another – When an atom absorbs energy, electrons are promoted to higher energy levels. – When the atom releases the absorbed energy, the electron falls back down to lower energy levels and EM radiation is given off (sometimes light) – These amounts are unique to every element

10. n=1 n=2 n=3 n=4 Spectrum UV IR VisibleVisible Ground State Excited State Excited State unstable and drops back down Energy released as a photon Frequency proportional to energy drop Excited State But only as far as n = 2 this time

11. Light and Energy Emission spectrum of each element is unique to that element

12. Electron Configurations Energy Levels – N= 1,2,3,4 Sub-level – Shapes – How many orbitals – How many electrons in each sub-shell

13. Electron Configurations

14. How electrons fill the orbitals is based on 3 rules: – The Aufbau principle – The Pauli Exclusion principle – Hund’s rule

15. Orbital Diagram

16. Electron Configurations The Pauli exclusion principle – Atomic orbital may describe at most 2 electrons – Electrons must have opposite spins – ↑↓

17. Electron Configurations Hund’s rule In a set of orbitals, the electrons will fill the orbitals in a way that would give the maximum number of parallel spins (maximum number of unpaired electrons) ↑ ↑ ↑

18. Writing Electron Configurations Full electron configuration Atoms and Ions – Al: 1s 2 2s 2 2p 6 3s 2 3p 1 – Al 3+: 1s 2 2s 2 2p 6 Orbital Configuration (arrows) Noble Gas Simplification for Atoms and Ion – [He] 2s 2 2p 6

19. Electron Configurations Identify – valence electrons – What element (based on electron configuration) 1s 2 2s 1 – Invalid electron configurations

20. Periodic Trends Groups Vertical column of the periodic table Indicates valence electrons Periods Horizontal row of the periodic table Period will indicate number of energy levels around the nucleus

21. Periodic Trends Atomic radius – increase as we go down a group – decrease we go across a period Ionic Radius – decreases when being converted to a positive ion – increases when being converted to a negative ion

22. Periodic Trends Electronegativity – decreases as you down a group – increases as you move across a period Ionization Energy – decreases as you go down a group – Ionization energy increases as you move across a period