November 06, Chapter 7 Atomic Struture. CHAPTER 7 Atomic Structure. Oct 27 9:34 AM ATOMIC STRUCTURE. Oct 27 9:34 AM

Transcription

1 CHAPTER 7 Atomic Structure ATOMIC STRUCTURE 1

2 The Wave Nature of Light Most subatomic particles behave as PARTICLES and obey the physics of waves. Visible light Ultravioletlight Wavelength Frequency (Hertz or s 1 ) (lambda) crest to crest or trough to trough (meters) (nu) the # of wave cycles that pass a given point. Wave motion: wave length and nodes 2

3 Electromagnetic Spectrum C = = 2.99 x 10 8 m/s speed of light Note that long wavelength > small frequency Short wavelength > high frequency Oct 6 2:14 PM Electromagnetic Radiation 3

4 Electromagnetic Radiation All radiation: λ. ν = c c = 2.99 x 10 8 m/s speed of light Indirect relationship Long wavelength = small frequency Short wavelength = high frequency Electromagnetic Spectrum Note the wavelength of visible light 4

5 Quantization of Energy Max Planck ( ) Solved the ultraviolet catastrophe The wave nature of light does not explain how an object can glow when its temperature increases. Quantization of Energy An object can gain or lose energy by absorbing or emitting radiant energy in QUANTA. Energy of radiation is proportional to frequency E = h ν h = Planck s constant = x J s 5

6 Photoelectric Effect Photoelectric Effect Classical theory said that E of ejected electron should increase with increase in light intensitynot observed! No e observed until light of a certain minimum E is used. Number of e ejected depends on light intensity. 6

7 Understand experimental observations if light consists of particles called PHOTONS of discrete energy. Therefore, if one knows the wavelength of light, one can calculate the energy in one photon, or packet, of that light: c = λν E = hν Energy of Radiation Calculate the Energy of 1.00 mol of photons of red light. 7

8 Another mystery involved the emission spectra observed from energy emitted by atoms and molecules. Classical Theory any amount of E could be gained or lost by an object this would result in a continuous line spectrum. NOT SO! When atoms or molecules change states, it absorbs or emits an amount of Energy Quantum Theory Ground state the state of lowest energy Excited state any state of higher E than ground state excited state Absorbs Emits Heat energy Light energy ground state Atomic Emission Spectrum produced when an e drops from higher energy to lower energy (fingerprints of elements) When absorption or emission from one energy state to another takes place, light with a certain is given off. Photon Streams of light particles (A quanta of light) Quantum Packet of energy If e s are in quantized energy states, then E of states can have only certain values. This explain sharp line spectra. Emission Spectrum 8

9 9

10 Excited atoms emit light of only certain wavelengths The wavelengths of emitted light depend on the element. Excited atoms can emit light. Here the solution in a pickle is excited electrically. The Na + ions in the pickle juice give off light characteristic of that element. 10

11 The energy absorbed or emitted from the process of electron promotion or demotion can be calculated by the equation: where R H is the Rydberg constant, J, and n i and n f are the initial and final energy levels of the electron. Note that emitting energy is exothermic 11

12 The Wave Nature of Matter If light can have material properties, matter should exhibit wave properties. For light: E = mc 2 E = hν = hc / λ Therefore, mc= h / λ For particles (mass)(velocity) = h / λ 12

13 Baseball (115 g) at 100 mph λ = 1.3 x cm e with velocity = 1.9 x 10 8 cm/sec λ = nm Uncertainty Principle W. Heisenberg Heisenberg showed that the more precisely the momentum (m x v)of a particle is known, the less precisely is its position known: In many cases, our uncertainty of the whereabouts of an electron is greater than the size of the atom itself! 13

14 Quantum or Wave Mechanics E. Schrodinger Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated. It is known as quantum mechanics. Quantum Numbers Solving the wave equation gives a set of wave functions, or orbitals, and their corresponding energies. Each orbital describes a spatial distribution of electron density. An orbital is described by a set of three quantum numbers. Nov 3 1:38 PM 14

15 Principle Quantum Number,n tells how far an e is from the nucleus Energy Level The larger the n, the further away the electron is from the nucleus, creating less attraction forces toward the nucleus. Outside e are likely to be the "bonding" e called valence e. n are integers 0. Principle quantum Number Orbital Quantum Number, Sublevels, l Describes the shape of the orbital (Area where e might be) Allowed values of l are integers ranging from 0 to n 1 & tells the number of nodal planes in the orbital Orbital Quantum Number 15

16 Orientation Quantum Number, m l Describes how the orbital is arranged/positioned in space s p d f Values are integers ranging from l to l: l,...+1,0, 1,... l orientations ml l Orbitals with the same value of n form a shell. Different orbital types within a shell are subshells. Orientation Quantum Number Electron Spin Magnetic Quantum Number,m s Describes the direction the e is spinning In the 1920s, it was discovered that two electrons in the same orbital do not have exactly the same energy. The spin of an electron describes its magnetic field, which affects its energy. Pauli Exclusion Principle Only 2 e can have the same set of 4 quantum numbers, they must have opposite spins paramagnetic ferromagnetic diamagnetic Spin Quantum Number 16

17 17

18 18

19 19

20 20

21 21

22 Energies of Orbitals For a one electron hydrogen atom, orbitals on the same energy level have the same energy. That is, they are degenerate. As the number of electrons increases, though, so does the repulsion between them. Therefore, in many electron atoms, orbitals on the same energy level are no longer degenerate. Nov 3 9:16 PM 22

23 Attachments 07M08AN2.avi