Atoms, Electrons and Light MS. MOORE CHEMISTRY

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1 Atoms, Electrons and Light MS. MOORE CHEMISTRY

2 Atoms Remember Rutherford??? What did he discover with his gold foil experiment. A: Atoms contain a dense nucleus where the protons and neutrons reside. ATOMS ARE MOSTLY EMPTY SPACE. Where do the electrons reside? A: around the outside.. But how does this work? iew=detail&mid=3b120fad330c7fcae2c83b120fad330c7fcae2c8&for M=VRDGAR

3 Light Energy in the form of a wave Peaks and Troughs Where do we see waves? Particle? Photon packet of energy

4 Frequency/Wavelength Frequency (ν) - # of wave peaks that pass a point per unit time (typically seconds) Wavelength (λ) distance between two peaks on a wave λ # per second = ν νλ = c Speed of Light (c) 3.00 x 10 8 m/sec

5 Light Electromagnetic Energy (Radiation) X-Rays Visible Light Infrared Microwaves Increasing λ Increasing ν & E Gamma X-Ray UV Visible Infrared Microwaves Radio/TV nm 1 mm 1 m

6 Energy Emission of Energy Atoms absorb energy Released at some random point later Specific frequencies of light released We see Light Energy Released Excited State New Excited State Energy Ground State

7 Atomic Line Spectra White Light is composed of all frequencies Excited atoms give off discreet frequencies of light Sodium Yellow Hydrogen Blue Neon Red Mercury Blue Lithium Red Potassium - Purple Atomic Line Spectra Lines of specific frequencies of light given off by an atom

8 Photons Photoelectric effect- is the emission of electrons from a metal when light shines on the metal. Max Planck in the 1900 s studied emission of light by hot objects. Quantum of energy- is the minimum quantity of energy that can be lost or gained by an atom. Einstein- proposed that electromagnetic radiation has a dual wave-particle. Photon- is a particle of electromagnetic radiation having zero mass and carrying quantum of energy.

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10 Energy Ground state- the lowest energy level of an atom. Excited state- a state in which an atom has a higher potential energy that it has in its ground state. When an excited atom returns to its ground state, it gives off the energy it gained in the form of electromagnetic radiation

11 Neil Bohr- Bohr s Model Neil Bohr- physicist proposed a hydrogen-atom model that linked the atom s electron to photon emission. -Energy levels represented as planetary orbits - Orbit closest to the nucleus = lowest energy state. n represents energy level.

12 Energy Excitation Hydrogen Atom 1 electron system Available energy levels Quantized (Discreet) Principle Energy Levels (n) Correspond to the distance from the nucleus to the electron Balmer series= 2 nd energy level. n=4 n=3 n=2 n=1

13 Emission/ Absorption - Emission- when an electron falls from to a lower energy level a photon is released. - - will see light - Absorption- energy must be added to an atom in order to move an electron from a lower energy level to a higher energy level.

14 Wave Mechanics Louis De Broglie Suggested that the electron could be treated as a wave Erwin Schroedinger Used the wave mechanical model to predict the structure of the atom Werner Heisenberg Showed that you cannot determine the velocity and the position of the electron Heisenberg Uncertainty Principle

15 Quantum Numbers n= quantum number region of space where electron may be found L= angular momentum Sublevel in which a group of atoms may be found m l = the magnetic quantum number Orientation of individual orbitals that may contain a pair of electrons M s = the spin quantum number. It can either spin up or down.

16 SubLevels l = 0 s sublevel l = 1 p sublevel l = 2 d sublevel l = 3 f sublevel

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18 Hunds Rule in a set of orbitals, the electrons will fill the orbitals in a way that would give the maximum number of parallel spins.

19 Aufbaus Principle Electrons occupy orbitals of lower energy first.

20 Pauli exclusion Principle No two electrons in the same atom can have the same set of four quantum numbers. Each orbital can only hold two electrons. In order for two electrons to occupy the same orbital they must have opposite spin.

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22 Electron Configuration Each shell has additional sublevel Important sublevels (s, p, d, f) Sublevels s: 1 orbital p: 3 orbitals d: 5 orbitals f: 7 orbitals Orbitals Each contain a maximum of 2 e -

23 Blocks in the Periodic Table

24 Energy Energy Order 5s 4d 4p 4s 3d 3p 3s 2p 2s 1s

25 Order of filling (Aufbau Chart) 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p 8s

26 Electron Configuration H: 1s 1 He: 1s 2 Li: 1s 2 2s 1 Be: 1s 2 2s 2 B: 1s 2 2s 2 2p 1 O 1s 2 2s 2 2p 4 Ne 1s 2 2s 2 2p 6

27 Practice writing electron configurations What is the electron configuration for Aluminum?

28 Noble Gas Electron Configuration Noble Gas Electron configuration contains a completely filled outer shell - shortcut is to use the noble gas associated with each element Ex: Na: [Ne]3s 1 Cr [Ar] 4s 2 3d 4 Practice:

29 Electron Configurations for charged atoms # Electrons= # protons If the Charge is positive subtract from the number of electrons Al 3+ = Na 1+ If the charge is negative add to the number of electrons S -2 = Highest energy level= outermost shell Add or subtract electrons from the outermost shell Ex: S -2 Ex: Fe 2+

30 Electron Configuration Cont Atoms that have the same electron configurations are called isoelectric(having the same # of electrons) Practice Write the electron configuration for Cu 2+ Write the electron configuration for P -3 What is its isoelectric pair?

31 Valence Electrons Electrons in the outermost shell or energy level n= energy level or shell The maximum number of electrons in the outermost shell is 8 Everything wants to have 8 Electrons= Octet Rule O:[He] 2p 2 2p 4 = Sn: Practice

32 Valence Electrons Cont: Periodic Table

33 Practice Identify the correct number of valence electrons for the following elements: Na: Au Cl -1 Fe +3

34 Completing the Octet Rule 1)identify how many valence electrons it has 2) determine how many electrons it needs to complete the octet rule Elements want to react with elements that will help to fill their outer shell (complete the octet rule) Argon Potassium Bromine Practice

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