Chapter Seven. The Electronic Structure of Atoms

Transcription

1 Chapter Seven The Electronic Structure of Atoms 1

2 Wave Theory Wave Repeating disturbance spreading out from a defined origin Characterized by wavelength, frequency and amplitude 2 Wavelength () Distance between identical pts Units some form of meters Frequency () Number of waves that pass through a point in 1 second Units of cycles/sec or Hz Amplitude Height of wave from center pt Intensity of wave

3 Electromagnetic Radiation 3 Electromagnetic Radiation Emission/transmission of energy in the form of waves with both electrical and magnetic components Travels at the speed of light, c c= 3.00 x 10 8 m/s Frequency & wavelength linked c=

4 What is the wavelength of an FM-radiowave with a 94.9 MHz frequency? 4 c = =3.00 x 10 8 m/s = c/ 94.9M Hz = 94.9 x 10 6 Hz = 94.9 x 10 6 /s 3.00 x10 s 8 m x s x m

5 5 Quantum Theory and the Photoelectric Effect

6 Max Planck s Quantum Theory Measurements show amount of energy emitted by an object at a certain temperature is directly related to its wavelength 6 Theorized that the energy must be in discrete amounts. May be in multiples of these discrete amounts E = h, E = 2h, E = 3h Called the smallest amount of energy a Quantum: Planck s constant: h = 6.63 X J s Didn t know why, but math worked over entire spectrum

7 Einstein and the Photoelectric Effect Experiment to prove why E= h Full spectrum light hits metal surface Energy transferred to electrons in metal Electrons break free and escape to anode Flow of electrons recorded with voltmeter 7 Results Light energy must have a certain threshold frequency to dislodge electrons Light energy has wave properties: E = h Light energy as a particle: E= KE electron + w (work needed to dislodge electrons) These particles of light later called photons

8 Practical Examples of the Photoelectric Effect What is the energy of a radiowave with a frequency of 94.9 MHz? x10 Js 94.9x10 26 E photon h x 6.23x10 J 1 s What is the energy per photon and per mole of photons of violet light, with a wavelength of 415 nm? E photon 4.79x10 9 hc.626x x10 photon Js 3.00x10 x s J m 1 x 415x x10 photons x 1mol m 2.88x10 mol 19 5 E mol J J What wavelength has an energy of E = 1.00 x J? x x Js 3.00x10 x s m 1 x 1.00x10 J m

9 9 Bohr s Theory of the Hydrogen Atom

10 Emission Spectra Pattern of radiation that is emitted when photons are removed from a substance Procedure Add energy to a substance Photons are emitted as a beam of light Separate wavelengths through a prism Record pattern on a photographic plate 10

11 Continuous vs. Line Spectra Continuous spectrum: Occurs when all visible light is present: white light 11 Line Spectrum Occurs when light is produced through an element Shows a pattern of lines characteristic of that element Can be used for identification

12 Elemental Line Spectra 12

13 Bohr s Hydrogen Atom Niels Bohr (1913): Electron energy (E n ) was quantized Only certain specified values allowed Stable levels called energy levels Photon moves from 1 level to another 13 The energy of each stable orbit E n = R H /n 2 n is the quantum number of the level integers only, 1,2 3.. Proportionality constant R H Rydberg constant R H = 2.18 X J

14 Energy Level Calculations All calculations done by comparing energy levels Electron moves between levels E =R H (1/n i2 1/n f2 ) 14 Energy emitted or absorbed High to low level energy released (-) Low to high level energy absorbed (+) Ground state: The lowest possible energy level Excited state: All other levels

15 Calculate the wavelength of the electron shift from n = 4 to n =2. Is light emitted or absorbed? 15 E 2.18 E x RH 2 2 ni n f 1 1 Jx x10 19 J E hc so hc E x10 Js 3.00 x10 m 1 7 x x 4.86 x10 19 s 4.09 x10 J m = 486 nm Visible blue green light is emitted (higher to lower n)

16 Quantum Numbers 16

17 Quantum Numbers and Atomic Orbitals 17 Atomic orbital A region in space with a high probability of finding an electron. Identified by 4 quantum numbers. Four Quantum Numbers 1. The principal quantum number (n) 2. The angular momentum quantum number (l) 3. The magnetic quantum number (m l ) 4. The electron spin quantum number (m s )

18 The Principal Quantum Number (n) Restricted to the positive integers: 1, 2, 3, 4, 5, 6 or 7 Indicates the shell or level of the orbital 18 Indicates the size of the orbital Integers correspond to row numbers in periodic table n=1 n=2 n=3

19 19 The Angular Momentum Quantum Number (l) Indicates orbital shape Designation: s, p, d or f Designates the subshell Values range from 0 to n-1 level Name s p d f l= 0 are called s orbitals Spherical l = 1 are called p orbitals 2 teardrops joined at center I= 2 are d orbitals l= 3 are f orbitals.

20 The Magnetic Quantum Number (m l ): Determines the orientation in space of the orbitals Integers from -l to +l 20 Determines the number of orbitals in a subshell The number of possible values for m l = 2l + 1 l = 1 # values = 3 Range: -1 to 1 m l

21 Electron Spin Quantum Number (m s ) 21 A magnetic field is induced by the moving electric charge of the electron as it spins Spins cancel one another No net magnetic field for the pair Allows 2 electrons to occupy 1 orbital Designation two values: +1/2 and 1/2

22 Quantum Numbers Summary 22

23 Electron Configurations 23 The energy of an electron is defined by both n and l Principle shells n = 1,2, 3, 4 or 5 Subshells l = 0, 1, 2, or 3 depends on n s p d orbitals present (f subshell not shown)

24 Orbital Filling in Multi-electron Atoms Fill low to high with 2 electrons per orbital Use chart to account for overlap of n values 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s Format: nl #e- 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 24

25 Electronic Configuration Defines the orbital for each electron # electrons = atomic number (Z) of atom Orbital Diagram Show boxes or lines for levels Place designate electrons 2 electrons allowed per orbital spdf Notation n designates subshell Letter designates l Use superscript for # electrons 25 Ne: Z = 10 1s 2 2s 2 2p 6 Na: Z = 11 1s 2 2s 2 2p 6 3s 1 or Na: [Ne]3s 1

26 Hund s Rule and the Aufbau Principle Hund's rule: Electrons in the same subshell occupy degenerate orbitals singly, before pairing Oxygen, O Z = s 2s 2 px 2 py 2 pz 26 The Aufbau Principle: Electron configuration of each element is based on the electron configuration of the element just before it in the periodic table. (Z = 1) H 1s 1 (Z = 2) He 1s 2 (Z = 3) Li 1s 2 2s 1

27 Magnetism in Multi-electron Atoms +1/2 & -1/2 spins will cancel if electrons paired No magnetic properties without spin present # unpaired electrons proportional to magnetic properties 27 Diamagnetic All electrons paired Ne: 1s 2 2s 2 2p 6 Paramagnetic At least 1 unpaired electron Na: 1s 2 2s 2 2p 6 3s 1

28 Exceptions To The Aufbau Principle Elements will fill out a lower energy subshell when possible Cr and Cu fill out their 3d shell before the 4s shell. Elements in the same family as Cr &Cu behave in same way. 28

29 Quantum Numbers and the Periodic Table Principle quantum number, n Row number of periodic table, values of 1-7 Angular momentum quantum number, l Specific area of periodic table, spdf Number of orbitals and electrons, m l & m s m l : Count blocks and divide by 2 m s : Count blocks 29

30 Electronic Configurations and the Periodic Table Add 1 electron for each block in the periodic table 30