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1 Electrons in Atoms Chapter 5.2-3. 2 Particle Nature of Light Max Planck (1858-1947) 1900 German physicist Studied light emitted from objects Found matter.

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Presentation on theme: "1 Electrons in Atoms Chapter 5.2-3. 2 Particle Nature of Light Max Planck (1858-1947) 1900 German physicist Studied light emitted from objects Found matter."— Presentation transcript:

1 1 Electrons in Atoms Chapter 5.2-3

2 2 Particle Nature of Light Max Planck (1858-1947) 1900 German physicist Studied light emitted from objects Found matter can gain/lose energy at small, specific amounts (QUANTA) Planck’s constant = 6.626 x 10 -34 J s

3 3 Planck’s work Increase radiation, increase frequency Violet light has more energy than red light Increase in whole # intervals 1hv, 2hv, 3hv, etc.

4 4 Photoelectric effect Photoelectron emitted from metal surface when light of certain frequency shines on surface Ex: calculator with photoelectric cells Light  electrical

5 5 Photoelectric effect is fun! with modification

6 6 Albert Einstein 1905 Photon  bundles of energy Particles of EM radiation with no mass that carry quantum of energy E photon = hv Photon must have minimum amount of energy (quanta) to eject photoelectron

7 7 Atomic Emission Spectra Set of frequencies of EM waves emitted by atom Absorb energy and become excited Release energy by emitting light Spectroscope prism inside that splits white light

8 8 Continuous Spectrum Continuum of visible colors from red to violet Look through spectroscope to white light What do you see?

9 9 Absorption Spectrum Continuous spectrum containing black lines at wavelengths associated with atom’s energy absorptions The black you see in the spectroscope is absorption

10 10 Emission Spectrum Colored lines associated with atom’s energy-level transitions Light reflected is the color we see Ex: leaves appear green b/c reflect green light and absorb the rest Lines of colors unique to element/color

11 11 5.2 Quantum Theory & the atom

12 12 Bohr Model of the Atom Described the energy states of H Quantized energy (only certain values allowed) Established Ground State Lowest allowable energy state 1 e-, many different excited states

13 13 e- moves in certain path Smaller path, lower atom’s energy state (energy level) Assigned quantum # (n) to each orbit n=1 Closest to nucleus

14 14 Faults to Bohr model Only works with H spectrum Not fully account for chemical behavior of atoms e- do not move around nucleus in circular orbits

15 15 Series Info Balmer series Series of visible lines based on e- dropping from higher-energy levels Lyman series (UV) e- drops into n=1 Paschen series (IR) e- drops into n=3

16 16 Quantum Mechanical Model of Atom 1926 Schrodinger derived an equation that led to the quantum mechanical model of the atom This is the currently accepted model

17 17 Heisenberg Uncertainty Principle fundamentally impossible to know precisely both the velocity and position of a particle at the same time because your method of determination will affect the particle.

18 18 5.3 Objectives Apply the Pauli exclusion principle, the aufbau principle and Hund’s rule to write e- configurations using orbital diagrams and e- configuration notation Define valence e- and draw e- dot structures representing an atom’s valence e-

19 19 5.3 e- configuration the arrangement of e- in an atom

20 20 Electrons fill energy levels according to 3 rules: 1. aufbau principle e- fills the lowest energy orbital available (e- are lazy) All orbits in a sublevel have the same energy Sublevels in an energy level have different energies (least s to greatest f) Sublevels from different energy levels may overlap

21 21 Electrons are LAZY!

22 22 continued 2. Hund’s Rule single e- with the same spin must occupy each equal-energy orbital before 2 e- can be in the same orbital e- are anti-social – 1 person to a seat on the bus before 2 will sit together

23 23 Electrons are ANTI- SOCIAL!

24 24 continued 3. Pauli Exclusion Principle two e- may be in the same orbital, but only if they have opposite spins

25 25 Electrons attract OPPOSITES!

26 26 Orbital Diagrams An empty box represents an unoccupied orbital Each box is labeled with principle quantum number and sublevel Fill boxes according to 3 rules until correct # of e- for the element are reached

27 27 e- config notation Principle energy level, sublevel, # of e- as subscript Label s, p, d, f on PT Move He adjacent to H for this purpose Start at beginning & write as you complete each sublevel d sublevel is on previous row’s energy level (row -1 for energy level) The f sublevel is (row -2 for energy level) In the transition elements (d box) a half filled or fully filled d sublevel is more important to stability than fully filled s sublevel Instead of s 2 d 4, the correct config is s 1 d 5 Instead of s 2 d 9, the correct config is s 1 d 10

28 28 Noble gas config Instead of writing out the complete config from 1s on, write the symbol for the noble gas of the previous row in brackets complete the config as normal

29 29 Valence e- Electrons found in the outer energy level determine the chemical property of the element group # for the representative elements represents the # of valence e- He is full at 2 e-, just as the rest are stable at 8 The p sublevel being full is the most stable config. d is next, then ½ d.

30 30 e- dot diagrams Write the symbol for the element pair the valence e- on each side of the symbol

31 31 Credits Information obtained from Glencoe Chemistry Matters Arranged by Michelle Estrada

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