Quantum Theory and Electron Configurations

Transcription

1 Chapter 5

2 Chapter 5

3 Quantum Theory and Electron Configurations

4 It s all about color In terms of atomic models, so far: Dalton (1803) = Tiny, solid particle Thomson (1897) = Plum Pudding model Electrons stuck on the outside of a big positive charge Rutherford (1911) = Positively-charged nucleus with electrons moving around it Rutherford s model of the atom not quite right Could not explain chemical properties of elements Could not explain color changes when metal is heated

5 Bohr Model of the Atom Niels Bohr s model of the atom Electron found only on specific, circular paths around nucleus Each orbit has fixed energy level Hypothesis: When electrons are excited (added energy), jump into higher energy levels. When they moved back into lower energy levels - gave off light. Electrons do not exist between levels (think of rungs on a ladder) Electrons absorb and emit only certain quanta (amounts) of energy Quantum of energy = fixed amount of energy required to move from one energy level to another energy level

6 Bohr s Model Nucleus Electron Orbit Energy Levels Chapter 5

7 Bohr s Planetary Model of the Atom Electrons must have enough energy to keep moving around the nucleus Electrons orbit nucleus in defined energy levels, just like planets orbit the sun Each energy level assigned a principal quantum number n. Lowest energy level called ground state (n=1) Higher energy levels (n=2, 3, 4...) excited states Model worked OK for hydrogen but not so good for other elements Nucleus n = 1 n = 2

8 Increasing energy Bohr s Model Fifth Fourth Third Second First Nucleus Chapter 5 Further away from the nucleus means more energy. There is no in between energy Energy Levels

9 Electron starts on lowest energy level (ground state) Add energy to electron moves to excited state Lowest energy level = ground state Higher energy levels = excited states Energy levels are not evenly spaced Energy Level 3 Energy Level 2 Energy Level 1 Nucleus

10 Electron starts on lowest energy level (ground state) Add energy to electron moves to excited state Energy Lowest energy level = ground state Higher energy levels = excited states Energy levels are not evenly spaced Electron returns to lower state emits/gives off quantum of energy Energy Level 3 Energy Level 2 Energy Level 1 Nucleus

11 Bohr used this theory to explain the lines in the atomic emission spectra for hydrogen Chapter 5

12 Each of these lines corresponds to different energy changes 434 nm 656 nm 410 nm 486 nm Chapter 5

13 Chapter 5

14 Chem I - Mon, 9/22/15 Do Now Get to work on the PEN worksheet from last class Homework MEAL paragraph if not finished Agenda History Intro to quantum Electron Config

15 Quantum-Mechanical Model of the Atom Since the Bohr model had a very limited use, a new and very different model of the atom exists The Quantum Mechanical Model (1926) contains: Quantum energy levels Dual wave/particle nature of electrons Electron clouds In the new model, don t know exactly where electrons are - only know probabilities of where they could be

16 Heisenberg Uncertainty Principle Heisenberg Uncertainty Principle = impossible to know both the velocity (or momentum) and position of an electron at the same time

17 Quantum-Mechanical Model of the Atom Orbital = region around nucleus where an electron with a given energy level will probably (90%) be found Four kinds of orbitals s - spherical in shape, lowest orbital for every energy level p - dumbbell shaped, second orbital d - complex flower shape, third orbital f - very complex shape, highest orbital

18 s-orbitals All s-orbitals are spherical. As n increases, the s-orbitals get larger.

19 p- orbitals Three p-orbitals: p x, p y, and p z Lie along the x-, y- and z- axes of a Cartesian system. Dumbbell shaped, gets larger as n increases

20 d and f - orbitals There are five d and seven f-orbitals.

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23 Quantum Mechanical Model Principle Energy Levels (n) Labeled from 1-7 First energy level is n=1 Contains sublevels (s, p, d and f) Each energy level contains the number of sublevels equal to its value for n If n=3, there are three sublevels

24 Chapter 5

25 Quantum Mechanical Model In each sublevel there are atomic orbitals Atomic orbitals describe a space where an electron is likely to be found Type of subshell Shape of orbitals Number of orbitals Orbital names s Spherical 1 s p Dumbbell 3 p x, p y, p z d Cloverleaf (and one donut) 5 f Multi-lobed 7

26 Quantum Mechanical Model Each orbital can contain two electrons. Since negative-negative repel, these electrons occupy the orbital with opposite spins.

27 Quantum Mechanical Model The total number of orbitals of an energy level is n 2. For the third principle energy level, n=3, which means there are 9 orbitals These orbitals are 3s, 3p x, 3p y, 3p z and the 5 d orbitals Remember, we no longer think of orbitals as concentric circles, but we can say that n=4 extends farther from the nucleus than n=1.

28 Valence Electrons Only those electrons in the highest principle energy level

29 Energ y Electron Configuration and Orbital Notation Aufbau Principle electrons fill lower energy orbitals first, bottom-up n=1 fills before n=3 Will an electron fill the 1s or the 2s orbital first? 2s 2p x 2p y 2p z 1s

30 Energ y 1s Electron Configuration &Orbital Notation Hund s Rule electrons enter same energy orbitals so that each orbital has one electron before doubling up Each of the first electrons to enter the equal energy orbitals must have the same spin If we have 7 electrons, how will they fill in the below orbitals? 2s 2p x 2p y 2p z

31 Energ y 1s Electron Configuration and Orbital Notation Pauli Exclusion Principle an orbital can contain no more than 2 electrons. Electrons in the same orbital must have different spins. If we have 8 electrons, how will they be arranged? 2s 2p x 2p y 2p z

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33 Apartment Analogy Atom is the building Floors are energy levels Rooms are orbitals Only two people per room

34 Orbital Diagrams Draw each orbital as a box. Each electron is represented using an arrow. Up arrows clockwise spin Down arrows counter-clockwise spin Determine the total number of electrons involved. Start with the lowest energy level (1s) and start filling in the boxes according the rules we just learned.

35 Energy Orbital Diagram 4p 3d 4s 3p 3s 2p 2s 1s

36 Increasing energy 7s 6s 5s 4s 3s 2s 7p 6p 5p 4p 3p 2p 6d 5d 4d 3d 5f 4f 1s Chapter 5

37 Increasing energy 7s 6s 5s 4s 3s 2s 7p 6p 5p 4p 3p 2p 6d 5d 4d 3d The first to electrons go into the 1s orbital Notice the opposite spins only 13 more 5f 4f 1s Chapter 5

38 Increasing energy 7s 6s 5s 4s 3s 2s 7p 6p 5p 4p 3p 2p 6d 5d 4d 3d The next electrons go into the 2s orbital only 11 more 5f 4f 1s Chapter 5

39 Increasing energy 7s 6s 5s 4s 3s 2s 7p 6p 5p 4p 3p 2p 6d 5d 4d 3d The next electrons go into the 2p orbital only 5 more 5f 4f 1s Chapter 5

40 Increasing energy 7s 6s 5s 4s 3s 2s 7p 6p 5p 4p 3p 2p 6d 5d 4d 3d The next electrons go into the 3s orbital only 3 more 5f 4f 1s Chapter 5

41 Increasing energy 7s 6s 5s 4s 3s 2s 1s 7p 6p 5p 4p 3p 2p 3d The last three electrons go into the 3p orbitals. Chapter 5 6d 5d 4d They each go into separate shapes 3 unpaired electrons 1s 2 2s 2 2p 6 3s 2 3p 3 5f 4f

42 Orbital Diagrams Orbital diagrams are used to show placement of electrons in orbitals. Need to follow three rules (Aufbau, Pauli, Hund s) to complete diagrams Li Be B C N Ne Na

43 Orbitals and Energy Levels Principal Energy Level Sublevels Orbitals n = 1 1s 1s (one) n = 2 2s, 2p 2s (one) + 2p (three) n = 3 3s, 3p, 3d 3s (one) + 3p (three) + 3d (five) n = 4 4s, 4p, 4d, 4f 4s (one) + 4p (three) + 4d (five) + 4f (seven) Chapter 5

44 shapes Summary Max electrons Starts at energy level s p d f Chapter 5

45 Increasing energy Orbitals and Energy Levels n = 4 and so on... 4f n = 2 n = 3 3d 3p 3s 4d 4p 4s 2p n = 1 2s 1s

46 Electron Configuration Let s determine the electron configuration for Phosphorus Need to account for 15 electrons Chapter 5

47 Writing Electron Configuration Determine the total number of electrons. Write the principle energy level number as a coefficient, the letter for the subshell, and an exponent to represent the number of electrons in the subshell. He: 1s 2

48 The Kernel (Noble Gas) Notation Determine the total number of electrons Find the previous noble gas and put its symbol in brackets Write the configuration from that noble gas forward as usual

49 Writing electron configurations Examples O 1s 2 2s 2 2p 4 Ti 1s 2 2s 2 2p 6 3s 2 3p 6 3d 2 4s 2 Br 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 5 Core format O [He] 2s 2 2p 4 Ti [Ar] 3d 2 4s 2 Br [Ar] 3d 10 4s 2 4p 5 Chapter 5

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